HSCChemistry

STATIC AND DYNAMIC EQUILIBRIUM

Inquiry question: What happens when chemical reactions do not go through to completion?

Students:

• conductpractical investigations to analyse the reversibility of chemical reactions, for example:
  • cobalt(II)chloride hydrated and dehydrated
  • iron(III)nitrate and potassium thiocyanate
  • burningmagnesium
  • burningsteel wool (ACSCH090)
• modelstatic and dynamic equilibrium and analyse the differences between open and closed systems (ACSCH079, ACSCH091)
• analyseexamples of non-equilibrium systems in terms of the effect of entropy and enthalpy, for example:
  
  • combustionreactions
  • photosynthesis
• investigatethe relationship between collision theory and reaction rate in order to analyse chemical equilibrium reactions (ACSCH070, ACSCH094)

FACTORS THAT AFFECT EQUILIBRIUM

Inquiry question: What factors affect equilibrium and how?

Students:

• investigate the effects of temperature, concentration, volume and/or pressure on a system at equilibrium and explain how Le Chatelier’s principle can be used to predict such effects, for example:

• heatingcobalt(II) chloride hydrate
• interactionbetween nitrogen dioxide and dinitrogen tetroxide
• iron(III)thiocyanate and varying concentration of ions (ACSCH095)

• explainthe overall observations about equilibrium in terms of the collision theory (ACSCH094)
• examinehow activation energy and heat of reaction affect the position of equilibrium

CALCULATING THE EQUILIBRIUM CONSTANT ( K_EQ )

Inquiry question: How can the position of equilibrium be described and what does the equilibrium constant represent?

Students:

• deducethe equilibrium expression (in terms of Keq) for homogeneous reactions occurring in solution (ACSCH079, ACSCH096)
• perform calculations to find the value of Keqand concentrations of substances within an equilibrium system, and use these values to make predictions on the direction in which a reaction may proceed (ACSCH096)
• qualitativelyanalyse the effect of temperature on the value of Keq (ACSCH093)
• conductan investigation to determine Keq of a chemical equilibrium system, for example:
  • Keqof the iron(III) thiocyanate equilibrium (ACSCH096)
• explorethe use of Keq for different types of chemical reactions, including but not limited to:
  • dissociationof ionic solutions
  • dissociationof acids and bases (ACSCH098, ACSCH099)

SOLUTION EQUILIBRIA

Inquiry question: How does solubility relate to chemical equilibrium?

Students:

• describeand analyse the processes involved in the dissolution of ionic compounds in water.
• investigatethe use of solubility equilibria by Aboriginal and Torres Strait Islander Peoples when removing toxicity from foods, for example:
  • toxinsin cycad fruit
• conductan investigation to determine solubility rules, and predict and analyse the composition of substances when two ionic solutions are mixed, for example:
  • potassiumchloride and silver nitrate
  • potassiumiodide and lead nitrate
  • sodiumsulfate and barium nitrate (ACSCH065)
• derive equilibrium expressions for saturated solutions in terms of K_spand calculate the solubility of an ionic substance from its K_sp value
• predictthe formation of a precipitate given the standard reference values for K_sp

REVISION – MODULE 4 – DRIVERS OF REACTIONS

ENTHAPLY CHANGE

Enthalpy (H) is a measure of the heat contentof a system. Absolute enthalpy cannot be measured. However, the change in enthalpy (∆H): the change in the heat content of system during a process, measured at constant pressure.

Δ𝐻 = 𝐻𝑝𝑟𝑜𝑑𝑢𝑐𝑡𝑠 − 𝐻𝑟𝑒𝑎𝑐𝑡𝑎𝑛𝑡𝑠

• 𝐻𝑝𝑟𝑜𝑑𝑢𝑐𝑡𝑠 < 𝐻𝑟𝑒𝑎𝑐𝑡𝑎𝑛𝑡𝑠, therefore Δ𝐻 is negative and heat energy released by the system (exothermic).
• 𝐻𝑝𝑟𝑜𝑑𝑢𝑐𝑡𝑠 > 𝐻𝑟𝑒𝑎𝑐𝑡𝑎𝑛𝑡𝑠, therefore Δ𝐻 is positive and heat energy absorbed by the system (endothermic).

ENERGY PROFILE DIAGRAMS

HESS’S LAW OF HEAT SUMMATION

‘The total enthalpy change in a chemical reaction is constant, whether the reaction is performed in one step or

several steps.’

Hess’s law is a form of the law of conversation of energy (First law of Thermodynamics).

BOND ENERGIES

Bond energy (or bond enthalpy) is the amount of energy required to break one mole of a bond in a gaseous molecule.

∆𝐻𝑟𝑥𝑛 = ∑ ∆𝐻𝑟𝑒𝑎𝑐𝑡𝑎𝑛𝑡 𝑏𝑜𝑛𝑑𝑠 𝑏𝑟𝑜𝑘𝑒𝑛 + ∑ ∆𝐻𝑝𝑟𝑜𝑑𝑢𝑐𝑡 𝑏𝑜𝑛𝑑𝑠 𝑓𝑜𝑟𝑚𝑒𝑑

∆𝐻𝑟𝑥𝑛 = ∑ ∆𝐻𝑏𝑜𝑛𝑑 𝑒𝑛𝑡ℎ𝑎𝑙𝑝𝑦 𝑜𝑓 𝑟𝑒𝑎𝑐𝑡𝑎𝑛𝑡𝑠 − ∑ ∆𝐻𝑏𝑜𝑛𝑑 𝑒𝑛𝑡ℎ𝑎𝑙𝑝𝑦 𝑜𝑓 𝑝𝑟𝑜𝑑𝑢𝑐𝑡𝑠

ENTROPY

Entropy (S) is a measure of how the available energy is distributed or dispersed amount particles in a system. It is also a measure of energy dispersal (function of temperature). Generally, low entropy à high entropy. (Chaos)

When energy can be distributed in more ways, there is a greater entropy

• Entropyis sometimes referred to as the measure of disorder or randomness.
• A system with greater possible arrangements (microstates Ω), or greater diversity of movement has higher

FACTORS THAT CHANGE ENTROPY

• Increasingthe number of particles à Increases microstates à Increase entropy
• Mixingdifferent types of particles à Increases microstates à Increase entropy
• Increasingthe volume of a container of gas à Increases microstates à Increases entropy

o The larger the volume, the more ways there are to distribute the energy

• Increasingthe number of particles in states with more freedom of movement (gas > liquid > solid)
• Moleculesbecoming more complex à Increases microstates à Increases entropy
• Increasetemperature à Increases microstates à Increases entropy

SECOND LAW OF THERMODYNAMICS

The second law of thermodynamics states that the entropy of the universe is always increasing.

Δ𝑆𝑢𝑛𝑖𝑣𝑒𝑟𝑠𝑒 = Δ𝑆𝑠𝑦𝑠𝑡𝑒𝑚 + Δ𝑆𝑠𝑢𝑟𝑟𝑜𝑢𝑛𝑑𝑖𝑛𝑔𝑠 > 0

CALCULATING ENTROPY IN CHEMICAL REACTIONS

Δ𝑆𝑟𝑥𝑛 = ∑ Δ𝑆𝑝𝑟𝑜𝑑𝑢𝑐𝑡𝑠 − ∑ Δ𝑆𝑟𝑒𝑎𝑐𝑡𝑎𝑛𝑡𝑠

GIBBS FREE ENERGY

In any process, the main form of interaction between the system and the surroundings is the exchange of heat.

• In an exothermic reaction, heat from the system enters the surrounding and increases temperature, whichwill increase its  The reverse will be true for an endothermic reaction.
• Ata lower temperature, the same amount of heat will cause a greater proportional change in 

The equation allows the comparison between the relative contributions of the two driving forces for a reaction, entropy and enthalpy.

- If 𝚫𝑮 < 𝟎 (Δ𝐻𝑠𝑦𝑠𝑡𝑒𝑚 < 𝑇Δ𝑆𝑠𝑦𝑠𝑡𝑒𝑚), a reaction is spontaneous
- If 𝚫𝑮 > 𝟎 (Δ𝐻𝑠𝑦𝑠𝑡𝑒𝑚 > 𝑇Δ𝑆𝑠𝑦𝑠𝑡𝑒𝑚), a reaction is non-spontaneous
-If 𝚫𝑮 = 𝟎, a reaction will occur both in the forward and reverse directions, equilibrium.

EQUATION SUMMARY

REVERSIBLE REACTIONS

OUTCOMES COVERED

• modelstatic and dynamic equilibrium and analyse the differences between open and closed systems (ACSCH079, ACSCH091)
• investigatethe relationship between collision theory and reaction rate in order to analyse chemical equilibrium reactions (ACSCH070, ACSCH094)
  
  • explainthe overall observations about equilibrium in terms of the collision theory (ACSCH094)

EQUILIBRIUM

For reversible reactions, a reversible arrow is used to indicate that both reactions are capable of proceeding.

Reactants ⇋ Products

• Physical changes are generally reversible

STATIC AND DYNAMIC EQUILIBRIUM

Reactions will proceed until either a static or dynamic equilibrium is reached. Equilibrium refers to the state of a

closed chemical system which:

• Theconcentrations of both reactant and products do not change with time
• Therate of the forward reaction is equal to the rate of the reverse reaction

Irreversible reactions (shown with a forward arrow →) that go to completion reach a static equilibrium.

Reversible reactions (shown with a reversible arrow ⇋) do not go to completion. In a closed system, reversible reactions will instead reach a state known as dynamic equilibrium.

• Atequilibrium, the rates of the forward and reverse reactions are the same, but non-zero.
• Theequilibrium is dynamic because there are changes occurring at the microscopic level, even though the system undergoes no change at the macroscopic level.

There are no macroscopic changes when a closed system is at equilibrium.

TYPES OF SYSTEMS

• Anopen system is a system where matter and energy can enter and leave.
• Aclosed system is a system where matter cannot enter and leave, but energy exchange can take place with the surrounding (in the form of pressure or heat)
• Anisolated system is a system where neither matter nor energy can 

RATES OF REACTION

1. Concentrationor Volume/Pressure
2. SurfaceArea
3. Temperature
4. Presenceof catalyst
5. Reactivity of reactants

COLLISION THEORY

In order for any reaction to proceed, reactants must collide. Particles need to collide with sufficient energy and in the correct orientation for it to be a successful reaction. A collision with sufficient energy and the correct orientation is called an effective collision.

The rate of reaction is how rapidly a reaction proceeds. The rate is defined as the change in the concentration of reactants or products over time. It is dependant on the frequency of effective collisions.

ADDITION/REMOVAL OF A REACTION COMPONENT

If the concentration of one reaction component increases, its rate of reaction will increase as there are more particles to collide with, thus increasing the frequency of effective collisions. The rate of that reaction will be relatively greater than that of the rate of the reverse reaction. This means that more products or reactants will being produced until equilibrium is reached.

Similarly, the reduction in one of the reaction components will reduce its rate of reaction. This occurs as there are fewer particles to collide with which reduces the frequency of effective collisions. The rate of that reaction will be relatively less than the rate of the reverse reaction. This means that more products or reactants will being produced until equilibrium is reached.

CHANGE IN VOLUME OR PRESSURE

A decrease in the volume of a chemical system involving gasses will result in gasses colliding more often. The gas particles will also be colliding with more energy as pressure is inversely proportional to volume. As particles are colliding more frequently and with more energy to overcome the activation energy barrier, the frequency of effective collision increases. Both the rate of the forward and reverse reaction will increase, however, the rate of reaction that uses the greatest number of moles will be relatively greater than the reverse reaction as there are more particles that can collide effectively with each other.

When the volume is increased and pressure is decreased, the partial pressures of all gasses will decrease. Both the rate of the forward and reverse reaction will decrease. The rate of reaction that produces more moles will be relatively greater than the reverse as the reaction is more likely to occur because it requires fewer particles to effectively collide.

Changing the overall pressure of a chemical system does not always cause a disturbance in equilibrium. For example, the addition of inert gasses.

CHANGE IN TEMPERATURE

By increasing the temperature of a chemical system at equilibrium, particles will possess more kinetic energy, which means that more particles (both reactants and products) have enough energy to collide and overcome the activation energy of the forward and reverse reactions.

For an exothermic reaction, the activation energy of the reverse reaction is higher than that of the forward reaction. An increase in temperature means that proportionally more products will be able to collide with enough energy in reverse reaction than the reactants. This causes the rate of the reverse reaction to occur at a faster rate than the forward reaction. Therefore, the concentrations of the reactants will increase whereas the concentration of the products will decrease until a new state of equilibrium is reached.

When the temperature is decreased, all the particles in the system lose energy which decreases both the rate of the forward and the reverse reaction. However, the rate of the reverse endothermic reaction will be relatively higher than the rate of the forward reaction.

OUTCOMES COVERED

• analyse examples of non-equilibrium systems in terms of the effect of entropy and enthalpy, forexample:

- combustionreactions 
- photosynthesis

PHOTOSYNTHESIS

Photosynthesis appears to be the reverse reaction of the combustion of glucose and may seem to be a reversible reaction. However, in nature, the process involves many individual irreversible steps which combine to give the overall reaction, hence photosynthesis is irreversible.

6CO_2(g) + 6H₂O_(l) → C₆H₁₂O_6(s) + 6O_2(g)

-∆𝐻 > 0 (endothermic)
-∆𝑆 > 0 (less moles, moving to more order)
-Non-spontaneous at all temperatures
-Chlorophyll is the catalyst for photosynthesis.
-UV rays drives the photosynthesis reaction.

SPONTANEITY AND EQUILIBRIUM

• -The Gibbs free energy change for this reaction is negative, therefore we would predict that the forward reaction is spontaneous.
  -The Gibbs free energy change for the reverse reaction will be the negative of this value, +33.3 𝑘𝐽𝑚𝑜𝑙⁻¹, so we would predict that the reverse reaction is nonspontaneous.
  -Entropy of mixing allows the reaction to be reversible.
  -The position with lowest free energy is somewhere in between pure reactants and pure products. This is the position of equilibrium.

The sign of ∆𝐺 indicates whether reactants or products will dominate the mixture with lowest free energy.

• -Since entropy of mixing always exist, technically no reaction is strictly irreversible.
  oHowever, if the position of equilibrium lies very close to the products, the reaction is called
  “irreversible” as the reverse reaction will not occur to any observable extent.
  ▪This is when Gibbs free energy is very negative.
  -Reversibility of a reaction can also be considered in terms of activation energy.
  oReactions are unlikely to be reversible as molecules will not have sufficient energy.
  oIn order for a reaction to be reversible, the forward and reverse reaction must have a small activation energy.

LE CHATELIER’S PRINCIPLE

OUTCOMES COVERED

• investigatethe effects of temperature, concentration, volume and/or pressure on a system at equilibrium and explain how Le Chatelier’s principle can be used to predict such effects, for example:
• heatingcobalt(II) chloride hydrate
• interactionbetween nitrogen dioxide and dinitrogen tetroxide
• iron(III)thiocyanate and varying concentration of ions (ACSCH095)
  • examinehow activation energy and heat of reaction affect the position of equilibrium

RATE OF REACTION (FROM MODULE 3)

• The rate of reaction is the speed with which reactants are converted to products, or how rapidly a reaction
• Rateis defined as the change in concentration of reactants or products over 
• Therate of reaction depends on the frequency of effective collisions.

FACTORS AFFECTING THE RATE OF REACTION

• Natureof reactants
• Concentration
• Surfacearea
• Temperature
• Catalysts
• Pressure/volume

NATURE OF REACTANTS

Every reaction has its own rate and its own activation energy, depending on the reactivity of the reactants. Aqueous solutions already have dissociated ions. They do not need to collide in any correct orientation and usually have very low E_a.

CONCENTRATION

The rate of reaction increases when the concentration of reactants is increased.

• Therate of reaction is directly proportional to the reactant 
• Increasingthe concentration increases the number of effective 
• Increasesthe number of particles in a given 

SURFACE AREA (PARTICLE SIZE)

The rate of reaction increases when the surface area of reactants is increased.

• Exposesmore particles to the  This increases the chance of a successful collision which therefore increases the rate of reaction.

TEMPERATURE

The rate of reaction increases when the temperature of the reactants is increased.

• Thetotal number of collisions 
• IncreasedKE, particles move  There is a greater chance of successful collision.
• When the temperature increases, the average kinetic energy of the molecules increase, thusmolecules move faster which means they collide more frequently.
• Theaverage energy of the collisions  Therefore, a higher fraction of collisions exceed activation energy.
• Ingeneral, reaction rate doubles every 10 degrees 

PRESENCE OF A CATALYST

A catalyst is a substance that increases the rate of reaction without being consumed. Catalysts work by allowing the reaction to take an alternative reaction pathway with a lower activation energy

WHAT IS LE CHATELIER’S PRINCIPLE

In 1888, a French chemist called Henri Le Chatelier (1850 1936) put forth the statement known as Le Chatelier’s.

Principle:

“If a system at dynamic equilibrium is disturbed by changing the conditions, the system undergoes a reaction which minimises the effect of the disturbance to attain a new equilibrium”

A chemical system at equilibrium can be disturbed in the following ways:

• Changein concentration
• Changein pressure à Change in volume àChange in concentration
• Changein temperature

Le Chatelier’s principle is a convenient method for predicting equilibrium shifts, but does not explain why it shifts. Collision theory explains the shift in equilibrium.

THE HABER PROCESS

CHANGE IN CONCENTRATION

N2_(g) + 3H2_(g) ⇋ 2NH3_(g)                    ∆𝐻 = −92𝑘𝐽 𝑚𝑜𝑙₋₁

ADDING A REACTION COMPONENT

Addition of H2(g) increases [H2(g)]. This will result in the rate of the forward reaction to increase, meaning the forward reaction has been favoured.

Since the rate of the forward reaction is different to the rate of the reverse reaction, the equilibrium has been disturbed.

Generally, the reaction that counteracts the disturbance will be favoured; in other words, its rate will increase relative to the other reaction.

REMOVING A REACTION COMPONENT

Removing N2(g) decreases [N2(g)]. This will result in the rate of the reverse reaction to increase, meaning the reverse reaction has been favoured.

The reaction that counteracts the disturbance will be favoured; in other words, its rate will increase relative to the other reaction.

CHANGE IN VOLUME (OR PRESSURE)

• MOLES, VOLUME AND PRESSURE

The pressure exerted by a gas arises from the force of the gas particles colliding with the walls of the container. Therefore, the pressure is proportional to the number of gas particles present.

Boyle’s Law states that pressure and volume are inversely proportional to each other.

CHANGE IN PRESSURE

Doubling the pressure of the container decreases volume. To counteract this effect, the reaction will shift in the direction that produces the least amount of moles. Therefore, the forward reaction will be favoured.

CHANGE IN TEMPERATURE

• EXOTHERMIC AND ENDOTHERMIC REACTIONS

The effect of a change in temperature on a reaction at equilibrium depends on whether the forward reaction is exothermic or endothermic.

If temperature of the system is increased, the reverse reaction will be favoured to counteract this effect. The reverse reaction is endothermic and will be favoured.

ADDITION OF A CATALYST

A catalyst increases the rate of a chemical reaction without being consumed, by providing an alternate pathway of lower activation energy. The addition of a catalyst reduces the activation energy of both the forward and reverse reaction by the same amount.

• Thereforethe addition of a catalyst will not disturb the equilibrium
• Theconcentrations of the components are not affected, by the system will reach equilibrium faster

ADDITION OF INERT GAS

The addition of an inert gas will increase pressure, but equilibrium will not be disturbed. This is because the concentrations of reactants and products remain the same, if the volume of the container does not change.

DIMERISATION OF NITROGEN DIOXIDE

When colourless dinitrogen tetroxide gas (N2O4) is enclosed in a vessel, a brown colour will appear indicating the formation of nitrogen dioxide (NO2). The intensity of the brown colour indicates the amount of nitrogen dioxide present in the vessel. The dimerization of nitrogen dioxide is an exothermic process

NOTES

• Ashift in the forward direction is called a shift towards the right 
• Ashift in the reverse direction is called a shift towards the left 
• Equilibriumwill shift to remove an added component (away from component)
• Equilibriumwill shift to replace a removed component (towards the component)
• Anincrease in volume will cause equilibrium to shift towards the side with more gas 
• Adecrease in volume will cause equilibrium to shift towards the side with fewer gas 

ANSWERING LE CHATELIER’S PRINCIPLE QUESTIONS

Clearly explain the effect of changes in conditions on the yield of equilibrium reactions.

• Statethat the change in reaction conditions disturbs the 
• “Accordingto Le Chatelier’s principle, the position of equilibrium shifts left/right”
• Justifythe shift:
  1. “Toreplace/remove”
  2. “Tothe side with more/less gas moles, to increase/reduce pressure”
  3. “Inthe exothermic/endothermic direction, to replace/remove heat”
•  … and minimise the disturbance
• Statethe effect of the shift, “Therefore, the concentrations of the reactants/products increase/decrease”

QUALITATIVE ANALYSIS OF EQUILIBRIUM

OUTCOMES

Investigatethe effects of temperature, concentration, volume and/or pressure on a system at
equilibrium and explain how Le Chatelier’s principle can be used to predict such effects, for example:

• heatingcobalt(II) chloride hydrate
• interactionbetween nitrogen dioxide and dinitrogen tetroxide
• iron(III)thiocyanate and varying concentration of ions (ACSCH095)

CONCENTRATION PROFILE DIAGRAMS

The changes occurring in a system can be identified by examining the shape of the line in a concentration profile diagram during a particular time period.

• Thex-axis is time or reaction progress
• They-axis is concentration or partial pressure

EQUILIBRIUM CONSTANT

OUTCOMES

• deduce the equilibrium expression (in terms of Keq) for homogeneous reactions occurring in solution(ACSCH079, ACSCH096)
• explorethe use of Keq for different types of chemical reactions, including but not limited to:
  • dissociationof ionic solutions

EQUILIBRIUM CONSTANT EXPRESSION

A quantitative way of describing the position of equilibrium is the equilibrium constant:

𝑎𝐴 + 𝑏𝐵 ⇋ 𝑐𝐶 + 𝑑𝐷

• Capitalletters represent chemical substances
• Lowercase letters represent the stoichiometric coefficients of the balanced equation

In an ideal system, the value of K is constant at constant temperature.

PURE, LIQUIDS AND SOLIDS

𝐶𝑎𝐶𝑂_3(𝑠) ⇋ 𝐶𝑎𝑂_(𝑠) + 𝐶𝑂2_(𝑔)

The concentrations of pure solids and pure liquids cannot change at a constant temperature.

• Becausethe equilibrium constant is only concerned with concentrations that change as they approach equilibrium, eliminate the terms for pure solids or liquids.
• Theconcentrations disappear from the equilibrium constant 
• Only gases and aqueous species appear in the equilibrium, except if all the reactant and products are allliquids.

OUTCOMES

• perform calculations to find the value of Keq and concentrations of substances within an equilibriumsystem, and use these values to make predictions on the direction in which a reaction may proceed (ACSCH096)

INTERPRETING THE EQUILIBRIUM CONSTANT

The larger the value of K, the further the equilibrium lies towards the RHS. A reaction with a very large K, proceeds almost to completion.

The smaller the value of K, the further the equilibrium lies towards the LHS. A reaction with a very small K, proceeds barely at all.

K AND THE DIRECTION OF REACTION

If concentrations are substituted into the expression at any point, the value is called Q, the reaction quotient.

The relative values of Q and K determines which way the reaction will proceed to reach equilibrium.

• If 𝑄 = 𝐾, then the system is at equilibrium
• If 𝑄 > 𝐾, then the backwards reaction will be favoured
• If 𝑄 < 𝐾, then the forwards reaction will be favoured

RICE CALCULATIONS

CALCULATIONS WITH QUADRATIC EQUATIONS

In some calculations, it will be required to solve a quadratic equation to determine 𝑥.

In these calculations, simplification can be used to make the calculation less complicated to solve:

• If K is very small, then 𝑥 will be very small, and the calculation can be simplified by using the following approximation: [𝑟𝑒𝑎𝑐𝑡𝑎𝑛𝑡]𝑖𝑛𝑖𝑡𝑖𝑎𝑙 − 𝑥 ≈ [𝑟𝑒𝑎𝑐𝑡𝑎𝑛𝑡]𝑖𝑛𝑖𝑡𝑖𝑎𝑙
• However, the assumption must be checked to be valid. The calculated value of 𝑥 must be less than 5% of the reactant’s initial concentration.

OUTCOMES

• qualitatively analyse the effect of temperature on the value of Keq (ACSCH093)

EFFECT OF TEMPERATURE ON K_eq

When temperature is increased, equilibrium shifts so that the endothermic reaction is favoured, whether it is the forward or the reverse reaction. Therefore, unlike changes in concentration and pressure, a change in temperature is the only factor that will change the value of K.

For exothermic reactions, K increases with lower temperatures and decreases with higher temperatures.

 For endothermic reactions, K decreases with lower temperatures and increases with higher temperatures.

IRON(III) THIOCYANATE

_OUTCOMES

• conductan investigation to determine K_eq of a chemical equilibrium system, for example:
  • K_eqof the iron(III) thiocyanate equilibrium (ACSCH096)

IRON (III) THIOCYANATE EQUILIBRIUM REACTION

When mixed together, aqueous solutions of iron(III) nitrate (Fe(NO3)3) and potassium thiocyanate (KSCN) combine, in a reversible exothermic process, to form the aqueous iron(III) thiocyanate complex ([Fe(SCN)]²⁺ ).

Iron(III) ions are very pale yellow colour, while iron(III) thiocyanate complex is an intense deep red colour. When diluted and at equilibrium, the colour of the mixture containing all three species is amber.

EQUILIBRIUM SHIFTS

A shift to the RHS, increase concentration of [Fe(SCN)]²⁺ → More intense

• Cooldown system: ice water bath
• Increase[Fe³⁺] and [SCN^-]: Add more FeCl3 and KSCN

A shift to the LHS, decreases concentration of [Fe(SCN)]²⁺ → Less intense

• Heatsystem: Hot water bath (Approx. 70 degrees Celsius)
• Decrease[Fe³⁺] and [SCN^-]: Add NaOH and AgNO3

NOTES

Since thiocyanate ions bind to iron via the nitrogen atom, the formula of the iron(III) thiocyanate complex is sometimes written as [Fe(NCS)]²⁺.

WHAT IS SPECTROPHOTOMETRY

Spectrophotometry is an analytical technique that is used to determine the concentration of a substance in a solution.

• Iron(III) thiocyanate strongly absorbs light ata wavelength of 447 nm and reflects light that is mostly orange-red, hence we see it as orange-red in 
• Iron(III)thiocyanate absorbs light blue 
• When it absorbs the light, electrons aretransitioning from lower to higher quantised energy levels. The energy gap between the levels is equal to the energy of the 447 nm wavelength 

A UV/Vis spectrophotometer like the one shown measure the intensity of light passing through a sample solution in a cuvette.

• Alamp provides white light (continuous source, which is narrowed and aligned into a beam using a slit.
• A prism splits the light into different wavelengths and is rotated so the desired wavelength of light passesthrough the monochromator (exit slit).
• Thelight beam passes through the sample, which absorbs a fraction of the  The greater the concentration of iron(III) thiocyanate, the greater the absorption of 447 nm light.
• The remaining light is transmitted through the sample and reaches a detector, which converts the amountof light to an electrical 

A spectrophotometer expresses the intensity of light in absorbance (A).

• A:Absorbance, a number typically between 3 – 2.5. It is dimensionless but often expressed in absorbance units (AU)
• I0:The intensity of light passing through the blank (reference) 
• I:The intensity of light passing through the analyte 

Absorbance is proportional to both concentration and the length of the sample, according to the Beer-Lambert Law.

•  𝗌: Extinction coefficient (also known as molar absorptivity), a constant that relates absorbance with concentration and path length, with the units 𝐿 𝑐𝑚−1 𝑚𝑜𝑙−1
• 𝒍: Path length of sample in cm
• 𝒄: Concentration of the substance in the sample in 𝑚𝑜𝑙 𝐿−1

SOLUBILITY

OUTCOMES

• describe and analyse the processes involved in the dissolution of ionic compounds in water

PROPERTIES OF IONIC COMPOUNDS

A process where equilibrium is often established is in the dissolution of ionic compounds in water.

• Ioniccompounds contain both cations and anions.
• Theyhave an electrostatic attraction between oppositely charged 
• Theyare:

- Brittle
- HighMP, BP
- Solidsat room temperature
- Moltenand aqueous state conducts electricity

SOLUTIONS OF IONIC COMPOUINDS

Soluble ionic compounds will dissolve in water to form aqueous solutions. When a soluble ionic compound is added to water, the ions at the surface of the crystal become surrounded by water molecules.

• Somemolecules of water separate from one another and are able to pack closer to the cations and anions of the ionic 

o They form ion-dipole forces between the molecules.

• Solute:Ionic salt
• Solvent:Water

When the ion-dipole forces between the ions and the permanent dipoles of the water molecules (adhesive forces) become stronger than the ionic bonds between the ions and the hydrogen bonding within the water (cohesive forces), the ions are dislodged from their position in the crystal.

• Theionic compound dissociates into its component 
• Theions become 

The solvated or hydrated ions are surrounded by a shell of water molecules known as a solvation layer.

• Thesolvation layer acts as a cushion and prevents a solvated anion from colliding directly with a solvated cation, and therefore keeps the ions in the 

Many ionic compounds are soluble in water and will dissociate to form aqueous solutions.

• Howevernot all ionic compounds are soluble in  For example, AgCl and Ca3(PO4)2 are insoluble.
• Theinsolubility of these ionic compounds is due to their strong ionic bonds which cannot be disrupted by the adhesive ion-dipole forces between solute ions and water molecules. Hence the ions do not become dislodged from their positions in the 

Entropy can also contribute to solubility. Most dissolutions are entropically favourable (∆𝑆 > 0), but some dissolutions are unfavourable.

• Thesolution consists of the solute and the solvent, so the change in entropy of each part can be considered separately to determine the overall change in entropy for the dissolution process:

∆𝑆_{𝑠𝑜𝑙𝑢𝑡𝑒} + ∆𝑆_{𝑠𝑜𝑙𝑣𝑒𝑛𝑡} = 𝑜𝑣𝑒𝑟𝑎𝑙𝑙 ∆𝑆

• Dependingon the relative magnitude of the effects, total entropy can increase or decrease for the overall system when dissolution occurs, although in majority of cases overall entropy will increase.
• An overall decrease in entropy sometimes occurs for ions with high charge density as they interact stronglywith water molecules, so the water molecules are held tightly around the solvated 

INCREASE IN ENTROPY

Solute in a solid state has a fixed ordered arrangement. Dissolved solute has free mobile ions. Therefore, there is an increase in entropy as the number of possible arrangement increases.

DECREASE IN ENTROPY

Pure water molecules are in random arrangements and are also mobile. When water becomes a solvent, molecules solvate the solute and has less possible arrangements decreasing entropy.

NOTES

“Clear” means that light can pass through the substance without being scattered.

“Colourless” means the substance is not coloured. Solutions are clear, but not necessarily colourless.”

OUTCOMES

conduct an investigation to determine solubility rules, and predict and analyse the composition of substances when two ionic solutions are mixed, for example:

• potassiumchloride and silver nitrate
• potassiumiodide and lead nitrate
• sodiumsulfate and barium nitrate (ACSCH065)

PRECIPITATION

The production of an insoluble compound, usually by reacting two soluble compounds.

                           NaCl_(aq) + AgNO3_(aq) → AgCl_(s) + NaNO3_(aq)

When two clear solutions are mixed together, an insoluble compound is formed. This is called a precipitate reaction. Precipitation reactions are used to remove minerals from drinking water, to remove heavy metals from wastewater and in purification plants of reservoirs.

SOLUBILITY RULES

OUTCOMES

• derive equilibrium expressions for saturated solutions in terms of Ksp and calculate the solubility of an ionic substance from its Ksp value.

SOLUBILITY

The solubility of a compound is the maximum amount of solute that can dissolve in a specific volume of solvent at a certain temperature.

• Solubilityis a physical constant, like melting point or boiling point

EQUILIBRIA IN SATURATED SOLUTIONS

When a solvent has dissolved all the solute it can at a given temperature, the resulting solution is saturated.

• Anysolution containing less solute than this is unsaturated.
• Ifmore solute is added to a saturated solution, no more will 

In a saturated solution, the system is in dynamic equilibrium:

• Allof the species – the ionic solid and dissolved ions – are present in the final 
• The forward reaction is occurring at the same rate as the reverse reaction. In other words, when an iondissolves, another ion is precipitating at the same time.

The equilibrium constant for these solution equilibria is called the solubility product constant (K_sp).

• Theequilibrium constant is for the equation written in the direction of the dissolution.

UNITS OF CONCENTRATION

MOLARITY

Molarity is the main unit of concentration used in chemistry:

-Thevolume of solution is expressed in litres (L).

•  The volume of the total solution (solute + solvent) is used calculate concentration.

PERCENTAGE BY MASS OR WEIGHT (% M/M OR % W/W)

• Percentagesby mass and weight are both extensively used in 
• % m/m is an abbreviation for the percentage mass of a substance relative to the total mass. This is thesame as percentage weight (% w/w).

PERCENTAGE BY VOLUME (% V/V)

MASS PER VOLUME (M/V)

• Massper volume is used in pharmacy and 
• Massper volume is used to measure the blood alcohol level of drivers.

-A blood alcohol level of 0.01 refers to 0.010/g / 100mL of blood

PARTS PER MILLION & PARTS PER BILLION

Partsper million and parts per billion (ppm and ppb) are useful when describing very dilute 

SUMMARY

OUTCOMES

• investigatethe use of solubility equilibria by Aboriginal and Torres Strait Islander Peoples when removing toxicity from foods, for example:
  • toxinsin cycad fruit

Many native foods eaten by Aboriginal and Torres Strait Islander people are poisonous and need to be detoxified before consumption

There are several physical and chemical processes user for detoxification:

COOKING OR ROASTING (CHEMICAL REACTION)

• Thefood items are heated in an oven or fire
• Theheat causes the toxins to decompose

LEACHING (PHYSICAL CHANGE)

• Thefood items are cut up into smaller pieces and soaked in running 
• Water-solubletoxins are washed 

FERMENTATION OR PROLONGED STORAGE (CHEMICAL REACTION)

• Thefood items are stored for long periods (several months to several years)
• Duringthis time, various biological processes occur that break down the 

They are digested by fungi, or broken down by the plant's natural enzymes. (Biological catalyst)

CYCADS (MACROZAMIA)

Cycads are palm-like plants that produce seeds in cones. Cycad seeds are a rich source of carbohydrates and have been eaten in regions of Northern Australia for thousands of years. However, the seeds contains highly toxic chemicals.

The two main types are cycasin and b-methylamino-l-alanine (BMAA).

To prepare them, Aboriginal and Torres Strait Islander people who ate these seeds would commonly prepare them by:

• Cookingthe seeds in an oven or fire

• Cuttingthe seeds open and leach them in running water for around a week
• Fermentationor storage for several months, or several years (some groups stored them for more than three years)

THE POISON, CYCASIN

• Cycasin(C8H16N2O7)
• Methylazoxmethanolglucoside
• MolarMass = 22g/mol
• pKa= 21 (measure of acidity)

SOLUBILITY

• 6g/L @ 25 in water/diluted ethanol
• Sparinglysoluble in absolute ethanol
• Insolublein benzene, chloroform + acetone

DECOMPOSITION

• Meltingpoint @144
• Decomposesat 154
• ProducesN2

MORETON BAY CHESTNUT (BLACK BEAN)

The Moreton Bay Chestnut is found on the east coast of Australia. It produces pods containing large seeds that are toxic. Eating unprocessed seeds causes vomiting and diarrhoea. When cooked, processed seeds taste like sweet chestnuts.

SOLUBILITY CALCULATIONS

OUTCOMES

• derive equilibrium expressions for saturated solutions in terms of Ksp and calculate the solubility ofan ionic substance from its K_sp value
• predictthe formation of a precipitate given the standard reference values for K_sp

CALCULATING 𝑲_𝒔𝒑 FROM SOLUBILITY

An equilibrium is only present in a saturated solution where the maximum amount of ionic compound has dissolved. Therefore, the solubility of a substance, whether given in moles per litre or mass per volume, can be used to calculate the solubility constant 𝐾_𝑠𝑝.

PREDICTING THE FORMATION OF A PRECIPITATE

Solubility constants can be used to predict if a precipitate will form when two solutions are mixed.

• The solubility constant for this reaction at 25℃ is 2.55 × 10−4
• This means [𝐵𝑎²⁺][𝑂𝐻⁻]² = 2.55 × 10−4 in a saturated solution.

If the concentration of Ba2+ and OH- are higher than the amounts in a saturated solution, 𝑄 > 𝐾_𝑠𝑝 and precipitation
will occur.

If the concentration of Ba²⁺ and OH^- are lower than the amounts in a saturated solution, 𝑄 < 𝐾𝑠𝑝 and precipitation will not occur.

THE COMMON ION EFFECT

In an aqueous solution of an ionic compound, the ions are dissociated.

• Thismeans that the ions separate into individual solvated 
• Ionsof the same species are indistinguishable, regardless of where they 

This means that in a saturated solution, if another substance is added that has an ion in common with the first substance, it will affect the position of equilibrium, leading to lower solubility. This is known as the common ion effect.

SOLUBILITY CURVES (SOLUBILITY AND TEMPERATURE)

Since dissolution is a reversible process, the position of equilibrium (extent of dissolution) will depend on temperature.

The relationship between solubility and temperature can be seen on a solubility curve, which shows the maximum amount of solute that can dissolve at a range of temperatures.

DISSOCIATION OF ACIDS

OUTCOMES

• explore the use of Keq for different types of chemical reactions, including but not limited to:

– dissociation of acids and bases (ACSCH098, ACSCH099)

COMMON ACIDS

ACID STRENGTH

Acids differ in their strength: the extent of ionisation or dissociation in water.

- A stronger acid will ionise further

Strong acids completely dissociate in water

- Straight arrows are used to indicate that these dissociations are irreversible and proceed to completion.

Weak acids partially dissociate in water.

- Reversible arrows are used indicate that these dissociations proceed to equilibrium.

The acid dissociation constant is the equilibrium constant for the dissociation (ionisation) of an acid into hydrogen

ions (H⁺) and an anion (K_a).

Acids that can produce more than one H+ ion are known as polyprotic acids.

Acid strength depends on the identity of the acid and the extent of its ionisation in water.

STRONG ACIDS

- Perchloric: HClO₄

- Hydriodic: HI

- Hydrobromic: HB_r

- Hydrochloric: HCl

- Nitric: HNO₃

- Sulfuric: H₂SO₄

PH SCALE

The acidity of a solution is determined by both the strength and the concentration of the acids present.

The concentration of the hydrogen ions in a solution are generally small. The pH scale, a logarithmic scale, is a

convenient way of expressing [H⁺] as a number generally between 0 (extremely acidic) and 14 (extremely basic).

A pH of 7 represents a neutral solution. The further away from 7, the more acidic or alkaline the solution.

pH can be calculated from the [H⁺] using the equation:

                                       pH = -[H+]

- p stands for −log₁₀, and the concentration of the hydrogen ions are in mol L⁻¹

- The notation pH derives from the French pouvoir hydrogene, meaning the “power of hydrogen”

DEGREE OF IONISATION

To calculate the percentage of any component in a sample, the formula is:

FOR A WEAK ACID HA: