module 5: equilibrium & acid reactions

UntitledModule 5: Equilibrium & Acid Reactions Static and Dynamic Equilibrium

Reversible Reactions
Reversible reactions are reactions that can occur in both directions allowing products to reform back into reactants.
Examples include:
iron (II) nitrate and potassium thiocyanate
Reverse reactions occur when the particles of products collide with more energy than the reverse activational energy and reform into reactant particles.
Reactions release/absorb small amount of energy into or from surroundings. In order to reverse a reaction, same amount of energy would need to be returned or exited from the system. Since that amount of energy is insignificant, it is easier to reverse the reaction.
Module 5: Equilibrium & Acid Reactions 2
Irreversible Reactions
Irreversible reactions are reactions that only occur in one direction not allowing products to reform back into reactants.
Examples include:
burning magnesium
burning steel wool
Irreversible reactions release/absorb large amounts of energy into or from surroudnings. Inorder to reverse the reactions, same amount of energy would need to be returned or exited from the system. Since that amount of energy is large, it is impossible to reverse the reaction.
Heating Cobalt (II) Chloride Hydrate
Cobalt (II) chloride hexahydrate turning into anhydrous cobalt (II) chloride demonstrates a reversible reaction.
When pink cobalt (II) chloride hexahydrate is heated causing water to evapourate it forms into blue anhydrous cobalt (II) chloride. Then adding water to blue anhydrous cobalt (II) chloride can turn it back into pink cobalt (II) chloride hexahydrate.
Iron (III) Nitrate & Potassium Thiocyanate
CoCl ⋅ 2 6H O 2 (s) CoCl + 2 (s) 6H O 2 (l)
Iron (II) nitrate reacting with potassium thiocyanate also demonstarted a reversible reaction
When iron (II) nitrate and potassium thiocyanate are mixed together they form iron(III) thiocyanate. This reaction is reversible so the reverse reaction occurs inorder for the system to reach equilibrium
Combustion Of Magneisum
The combustion of magnesium demonstartes a irreversible reaction
Magnesium can be burned in oxygen its froms a white powder called magneisum oxide. This reaction is irreversible since conbustion reactions release alot of heat energy which cannot be returned back into the system.
Combustion Of Steel Wool
The combustion of steel wool also demonstarted a irreversible reaction

model static and dynamic equilibrium and analyse the differences between open and closed systems
Thermodynamic Systems
A system is a part of a universe that is of interest to us like a particular chemical or physical reaction. While everything around the system - the rest of the universe - is regarded as the surroudnings.
Fe + 3+ (aq) SCN −
(aq) FeSCN 2+ (aq)
4Fe +(s) 3O (g)2 2Fe O 2 3 (s)
A system can either be open, closed or isolated.
Open System
In an open system both matter and energy can be interchanged between the system and its surroundings.
An example of a open system would be boiling water in a beaker. Heat energy is transferred into the beaker (system) while water vapour is exiting the beaker into the surroundings.
Closed System
In a closed system only energy can be exchanged between the system and its surroundings.
An example of a closed system would boiling water in a coverd beaker. While heat energy is still transferred inside the beaker, the water vapour is unable to escape into its surroundings since it is restricted by the cover.
Isolated System
In a isolated system, niether matter nor energy can be exchanged between the system and its surroundings.
The only true isolated system is the entire universe.
Equilibrium Systems
Reversible and irreversible reactions will eventually reach a state called chemical equilibrium. A chemical equilibrium occurs when the concentrations of
the products and reactants remain constant so that there is no observable change.
The two types of chemical equilibrium are dynamic equilibrium and static equilibrium.
Dynamic Equilibrium
Dynamic equilibrium is reached when the rate of the forward reaction is equal to the rate of the reverse reaction. Dynamic equilibrium can be only achieved by a reversible reaction in a closed system.
Even thought at a macroscopic level there is no observale change between the concentration of products and reactants, at a microscopic level the reactants are turning into products at the same speed as the products are turning into reactants.
Let us consider the system:
Now, take 2 100mL beakers labelled A and B. Fill A with 100 mL water and leave B empty. This represents initial concentrations of reactant A and B. Now to initiate the reaction, fill up a 10 mL pippette from A and transfer its contents into B. Take a 5 mL pipette and repeat procedure, transferring water from A to B with 10 mL pipette, then transfer back from B to A with 5 mL pipette.
Eventually the level of water in both beakers remains constant, which is the point where dynamic equilibrium has been reached.
Static Equilibrium
Static equilibrium is reached when the rate of the forward and reverse reaction is equal to zero. Static equilibrium can be achieved by reversible and irreversible reactions in both open and closed systems.
Unlike dynamic equilibrium, both at a macroscopic and microscopic level the reaction has completely stopped. This occurs when the activational energy of the forward and reverse reaction is so high that either reaction does not proceed.
A B

analyse examples of non-equilibrium systems in terms of the effect of entropy and enthalpy, for example: – combustion reactions – photosynthesis
Enthalpy
Enthalpy refers to the total heat content of a system. Reactions that decrease enthalpy are favourable while reaction that increase enthalpy are less favourable.
Enthalpy change is the total heat energy absorbed or release by a system during a chemical or physical reaction. The more negative the enthalpy change, the more spontaneous a reaction.
Entropy
Entropy refers to a system's thermal energy that is not available to do useful work. Reactions that increase entropy are favourable while reactions that decrease entropy are less favourable.
The more positive the entropy change, the more spontaneous the reaction.
Gibbs Free Energy
Gibbs free energy is a value that helps predict the spontaneity of a reaction. The spontaneity of a reaction is not related to the reaction's speed.
when:
Non - Equilibrium Systems
A non equilibrium system involves reaction that are irreversible and never reach an equilibrium. Examples include photosynthesis and combustion.
Combustion
A combustion reaction is always exothermic meaning they have a negative enthalpy change (favourable). They also increase entropy making them have a
ΔG = ΔH − TΔS
positive entropy change (favourable).
Due to the negative enthalpy change and positive entropy change, they are spontaneous and exergonic. Therefore, combustion reactions are highly favourable in the forward direction.
Example of a combustion reaction includes the combustion of glucose (cellular respiration)
Photosynthesis

investigate the relationship between collision theory and reaction rate in order to analyse chemical equilibrium reactions
Collision Theory
Collision theory states that chemical reactions occur as a result of collision between reacting particles. For a successful collision, particles must:
collide and make contact
collide with the correct orientation to form bonds
For the collision theory we assume:
particles size is neglible relative to distance between them
particles move randomly inside the container
collisions are elastic (no energy is lost)
Rate Of Reaction
C H O + 6 12 6 (s) 6O → 2 (g) 6CO + 2 (g) 6H O 2 (l)
6CO + 2 (g) 6H O → 2 (l) C H O + 6 12 6 (aq) 6O 2 (g)
The rate of reaction is the change of concentration over time - the speed of the reaction. Factors affecting the rate of reaction include:
surface area
Concentration
When concentration increases, so does the rate of reaction since there are now more particles which are more frequently colliding.
Temperature
When temperature increases, so does the rate of reaction because of the increased speed of particles which are more likely to collide with high energy.
Pressure
When pressure increases, the rate of reaction increases because particles come close toagther increasing chances of collisions.
Catalyst
A catalyst is a substance that speeds up the reaction rate without being used up by providing a alternate reaction pathway and lowering the activational energy.
Catalyst do not influence a reaction thermodynamiclaly (do not affect enthalpy change nor entropy change).
Equilibrium Reactions
Thermodynamic and kinetic play an important role in an equilibrium system.
The position of equilibrium is determined by thermodynamics. Whether the equilibrium is shifted left or right is dependant on how thermodynamically favourable the forward and reverse reactions are.
The rate of reaction is determined by kinetics - how fast a system will reach equilibrium.
Factors That Affect Equilibrium

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: – heating cobalt(II) chloride hydrate – interaction between nitrogen dioxide and dinitrogen tetroxide – iron(III) thiocyanate and varying concentration of ions
Le Chatelier's Principle
Le chatelier's principle states that if a stress is applied to a chemical system at equilibrium, the system will shift to minimize the change and establish a new equilibrium.
The shift refers to a change that temporary increases in reaction rate of the forward or reverse reaction. So, the reaction in one direction will predominate for a certain period until reaction rate equals and new equilibrium is achieved.
A stress applied could be:
Change in concentration
Dilution
Changing Concentration
If the concentration of a reactant is increased, or the concentration of a product is decreased, the system will shift forwards by using up reactants to make more products.
If the concentration of a product is increased, or the concentration of a reactant is decreases, the system will shift backwards by using up products to make more reactants.
The concentration of the species at the new equilibrium is not the same as the original concentrations at original equilibrium. Adding or removing pure solids or liquids does not change their concentration.
Changing Pressure
For a system's equilibrium position to shift when a change in pressure has occured there must have been:
a change in the partial pressure of the gases in the system
at least one gaseous reactant or product in the equation
different number of gaseous moles on either side of the equation
A change in pressure is caused when volume is changed since pressure is inversely proportional to volume. When pressure increases by decreasing volume, system favours the side with fewer gas moles to counter the increase in pressure. When pressure decreases by increasing volume, system favours the side with greater gas moles to counter decrease in pressure.
Total pressure can be increased by adding an inert gas which would not effect equilibrium position as partial pressure of reactants and products remain unchanged.
Dilution
Dilution is the process of adding water, increasing total volume. For an equilibrium system in solution, diluting is similar to increasing volume of a gaseous system. Diluting reduces the number of particles per unit volume (concentration) for all species.
This results the system to favour the side with a greater number of dissolved particles. Since dilution is a irreversible process the volume of an equilibrium system in solution cannot be decreased without removing some reactant / products.
Changing Temperature
When a system's temperature is altered, the system will shift in the direction that minimises the temperature change. Effect of temperature on an equilibrium system depends on whether the reaction is endothermic or exothermic.
Increasing temperature of a system also increases energy of substances. In order to counteract this change, heat energy must be used up by the system, meaning endothermic reaction is favoured.
Inversely, decreasing temperature of a system also decreases energy of substances. In order to counteract this change, heat energy must be produced, meaning exothermic reaction is favoured.
Presence Of Catalyst
A catalyst is a substance that increases the rate of a chemical reaction without getting consumed in the reaction or undergoing any permanent chemical changes itself.
When a catalyst is added to an equilibrium system, the rate of forward and reverse reaction is increased to the same extent. Since the relative rates of both reactions are unchanged, the equilibrium position will remain unaffected. A catalyst will only increase the rate at which a reaction reaches equilibrium.
Heating Cobalt (II) Chloride Hydrate
Recall the reversible reaction where pink cobalt (II) chloride hexahydrate could be heated to form blue anhydrous cobalt (II) chloride.
This reaction is endothermic. So if heat is added, the equilibrium position shifts right to favour forward endothermic reaction which absorbs heat. Inversely if heat is taken away, the equilibrium position shifts left to favour reverse exothermic reaction which releases heat.
Changing pressure would not affect the equilibrium position since no gasous reagents are present. And since both salts are pure solid, adding or removing either wouldn't change concentration and hence doesnt alter equilibrium position.
Nitrogen Dioxide & Dinitrogen Tetraoxide
When brown nitrogen dioxide gas or colourless dinitrogen tetraoxide gas is placed into a closed vessel, an equilibrium is established between the two:
This reaction is exothermic. So if temperature is increased, the equilibrium position shifts left to favour reverse endothermic reaction which absorbs heat. Inversely, if heat is taken away, the equilibrium position shifts right to favour forward exothermic reaction which releases heat.
If pressure were increased, the equilibrium shifts right to favour forward reaction with fewer gas moles to counteract pressure increase. Inversely, if pressure decreases, the equilibrium shifts left to favour reverse reaction with more gas moles.
If NO2 gas is added or N2O4 gas is removed, the equilibrium shifts right to favour forward reaction to use up NO2 and produce N2O4. Inversely, if N2O4 gas is added or NO2 gas is removed, the equilibrium shifts left to favour reverse reaction to use up N2O4 and produce NO2
Iron (III) Nitrate & Potassium Thiocyanate
Again recall the reversible reaction between yellow iron (III) nitrate & colourless potassium thiocyanate to produce red iron (III) thiocyanate.
This reaction is exothermic. So if temperature increases, the equilibrium position shifts left to favour reverse endothermic reaction which absorbs heat. Inversely,
CoCl ⋅ 2 H O 2 (s) CoCl + 2 (s) 6H O 2 (l) ΔH > 0
2NO 2 (g) N O 2 4 (g) ΔH < 0
Fe + 3+ (aq) SCN −
if temperature decreases, the equilibrium position shifts right to favour forward exothermic reaction which releases heat.
Changing pressure would not affect the equilibrium position since no gasous reagents are present
Adding Fe(NO3)3 will increase , causing equilibrium position shift right to favour forward reaction and use up excess iron (III) ions. Adding KSCN will increase , causing equilibrium position shift right to favour forward reaction and use up excess thiocyanate ions. Adding NaH2PO4 will decrease

explain the overall observations about equilibrium in terms of the collision theory & examine how activation energy and heat of reaction affect the position of equilibrium.
Collision Theory & Equilibrium
Le chatelier's principle aligns very well with the collision theory.
Concentration
If concentration of a reactant increases, more reactant molecules are present and moving randomly. As such, there is a higher liklehood of reactant molecules colliding with each other causing a reaction to take place.
As new equilibrium is apporached, there becomes fewer reactant molecules and more product molecules decreasing freqeuncy of reactant molecules colliding and increasing frequency of product molecules colliding until the rate of reaction of both forward and reverse reaction becomes equal and new equilibrium is reached.
Pressure
When pressure is increased by decreasing volume, particles are closer together and will find themselves at a higher concentration. Reaction rate of reverse and forward reaction will increase since collisions become more frequent. However, the reaction rate of the reaction that produces less moles of gas will be higher because a higher concentration has a greater effect on the side which has greater number of moles.
[Fe ]3+
This is why adding inert gases doesnt effect equalibrium position because they dont change concentrations of reactants and products even though total pressure increases.
Temperature
When temperature increases, molecules gain kinetic energy causing them to move more rapidly and with mroe energy. For a successful collision particles must collide and with sufficient energy.
Since alot more molecules now possess sufficient energy due to an increase in temperature, it will increase the likelihood of a collision being successful.
Calculating the Equilibrium Constant

deduce the equilibrium expression (in terms of Keq) for homogeneous reactions occurring in solution
Homogenous Reactions
A hemogenous reaction is any reaction which occurs in a single phase - all the products and reactants are of the same state (solids, liquids or gas)
Heterogenous Reactions
A heterogenous reaction is any reaction which does not occur in a single phase - products or reactants are of different states
A + (g) B → (g) C (g)
The Equilibrium Constant
The equilibrium constant Keq is a quantitative value that describes the position of equilibrium. The larger the value of Keq, the more equilirbium is shifted to the right. Inversly, the smaller the value of Keq, the more equilibrium is shifted to the left.
It is calculated as the product of the concentrations of each product to the power of its coefficient, divided by the product of the concentrations of each reactant to the power of its coefficient.
The equilibrium constant for the system:
Is given by:
assuming A, B, C and D are either gaseous or in solution. Solids and pure liquids are omitted from the equilibrium constant expression because their molar concentrations do not change and stay constant throughout the reaction. They are assigned a concentration value of 1 because their concentrations do not depend on how much of the pure substance is present.
This is because the concentration of a solid or pure liquid equal to density over molar mass. Since density is an intensive property – do not depend on the amount of matter present, density remains constant. Hence, solids and pure liquids are not included in the calculation of Keq.
Importance Of An Equation
The equilibrium constant depends on the chemical equation for a particular reaction.
If an equation is reversed, Keq is inverted
If coefficents are doubled, Keq is squared
If coefficents are halved, Keq is square rooted
Reaction Quotient
aA + bB cC + dD
eq b
The reaction quotient (Q) is a mathematical relationship between concentrations of the reactants and products in a chemical system similar to the equilibrium constant.
The reaction quotient for the system:
Is given by:
While the equilibrium constant only uses the concentration of the reactants and products when the system is at equiibrium, the reaction quotient uses the concentration of products and reactants at any point in time.
If Q < K, system equilibrium will shift to the right since they are not enough products.
If Q > K, system equilibrium will shift to the left since they are not enough reactants.

perform calculations to find the value of Keq and concentrations of substances within an equilibrium system, and use these values to make predictions on the direction in which a reaction may proceed
Determining Equilibrium Constant Expression
aA + bB cC + dD
[C] [D]c d
To determine an expression for the equilibrium constant of a system at equilibrium a full balanced chemical equation describing the reaction is needed.
1. For the system:
2. For the system:
3. For the system:
4. For the system:
Calculating Equilibrium Constant
K =eq [NO ]2 2 [N O ]2 2
NaCl (s) Na + + (aq) Cl − (aq)
K =eq [Na ][Cl ]+ −
(aq) 2H + (aq)
[CO ][H ]3 2+ + 2
CO + 2 (g) H O 2 (l) HCO + 3 − (aq) H + (aq)
K =eq [CO ]2
To calculate the equilibrium expression for a reaction, not only a full balanced chemical equation describing the reaction is needed but also the moles or concentrations of every specie in the reaction at equilibrium.
1. At equilibrium, concentration of ammonia is 0.02 mol/L, while the concentrarion of nitrogen is 0.0031 mol/L and the concentration of hydrogen is 0.012 mol/L. Calculate the equilibrium constant.
2. At equilibrium, a 500 mL aqueous solution contains 0.020 mol cobalt (II) hexahydrate ion, 0.40 mol chloride ion and 0.015 mol tetrachloridocobaltate ion.
N + 2 (g) 3H 2 (g) 2NH 3 (g)
[NH ] =3 0.02 mol/L
[N ] =2 0.0031 mol/L
[NH ]3 2
K =eq 7.5 × 10 (2 s.f.)−4
(Co(H O) ) + 2 6 2+
(aq) 4Cl − (aq) (CoCl) + 2−
(aq) 6H O 2 (l)
[Co(H O) ] =2 6 2+ =0.500 L
0.020 mol 0.040 mol/L
[CoCl ] =2− =0.500 L 0.015 mol 0.030 mol/L
K =eq [Co(H O) ] [Cl ]2 6 2+ − 4
[CoCl ]2−
Harder Calculations Involving Equilibrium Expression
Harder questions involving calculating equilibrium dont give the cocentrations of all species in the reaction but have to be calculated using a method called ICE
The steps to the ICE method are:
1. Determine the known quantities and convert to concentrations
2. Set up an I.C.E table with all known concentrations
3. Identify the specie for which initial and equilibrium concentration is known and calculate its change in concentration
4. Use stoichiometric ratios to find the change in concentration for all other species.
5. Use initial and change in concentration to calculate equilibrium concentration of all species.
6. Determine the equilibrium constant expression and substitute in the equilibrium concentration
1. A 10.0 L vessel intially contained 0.250 mol nitric oxide and 0.160 mol oxygen. After equilibrium, 0.100 mol nitric acid remained. Calculate the equilibrium constant.
K =eq 0.040 × 0.804 0.030
K =eq 1.8 (2 s.f.)
2NO + (g) O 2 (g) 2NO 2 (g)
[NO] =i =10.0 L 0.250 mol 0.0250 mol/L
[O ] =2 i =10.0 L 0.160 mol 0.0160 mol/L
[NO] =eq =10.0 L 0.100 mol 0.0100 mol/L

qualitatively analyse the effect of temperature on the value of Keq
Changes In The Equilibrium Constant
The value of the equilibrium constant is only altered by the temperature. If pressure or concentration of a species is changed, the value of the equilibrium constant will remain the same.
This is because, when pressure or concentration changes the system will shift but when a new equilibrium is established, the concentrations of each specie will be such that the equilibrium constant remains unchanged, if temperature remains constant.
Effect of a change in temperature on Keq depends whether the reaction is exothermic or endothermic
Effect Of Temperature On Keq (Exothermic Reaction)
If the temperature of a exothermic reaction increases then the value of Keq decreases. Inversely, if the temperature of a exothermic reaction decreases then the value of Keq increases.
Effect Of Temperature On Keq (Endothermic Reaction)
K =eq [NO] [O ]2 2
[NO ]2 2

conduct an investigation to determine Keq of a chemical equilibrium system, for example: – Keq of the iron(III) thiocyanate equilibrium
Investigating Keq in the iron(III) thiocyanate equilibrium
The system is:
In the system:
is colourless
is an intense red colour
To determine Keq, we can combine known amounts of Fe3+ and SCN-, then measure concentrations of iron thiocyanate using spectrophotometry at equilibrium.
The first step is to combine known amounts of iron (III) and thiocyanate ions. Once solutions have been combined, test tube is left until colour remains unchanged indicating equilibrium. Next step is to measire equilibrium concentrations of iron (III) thiocyanate using a spectrometer.
There is not universal mathematical conversion to determine absolute concentration. Hence, a calibration curve is required. This is done by taking standard solutions of known iron (III) thiocyanate concentrations and measureing their absorbances.
Spectrophotometer

explore the use of Keq for different types of chemical reactions, including but not limited to: – dissociation of ionic solutions – dissociation of acids and bases
Dissociation Of Ionic Solutions
The equilibirum constant can be used to describe the dissociation of ionic solutions. The equilibrium constant, written for the dissolution of a solid into pure water is knows as the solubility product Ksp.
For the system:
Its solubility product if given by:
The Ksp for a solid describes its solubility. The higher the Ksp, the more the equilibrium is shifted to the right and the soluble the substance is since at equilibrium more of the salt exists as dissociated, hydrated ions in solution.
Conversely, a low Ksp means that the salt is not very soluble, as the equilibrium is shifted left and most of the salt exists in solid precipitate form.
A B b a (s) aA + + (aq) bB − (aq)
K =sp [A ] [B ]+ eq a −
eq b
Dissociation Of Acids
The equilibrium constant can also be used to describe the dissociation of acids. When an acid ionises in water, it gives an to water which produces a hydronium ion ( ), conjugate base.
where:
is the conjugate base
The conjugate base is the acid molecules minus a . The degree to which this happens depends on the strength of the acid. The stronger the acid, the greater the degree of ionisation, and the more this equilibrium is shifted to the right.
The equilibrium constant for the ionisation of an acid is called the acid dissociation constant, .
The higher the value of , the more the equilibrium is shifted right, which means that more of the acid ionised, which describes a stronger acid. Thus, a higher means a stronger acid.
Dissociation Of Bases
The equilibirum constant can also be used to describe the dissociation of bases. When a base ionises in water, it takes a from water which produces hydroxide ions ( ) and a conjugate acid.
where:
is the conjugate acid
The conjugate acid is the base molecules plus an extra . The degree to which this happens depends on the strength of the base. The stronger the base, the greater the degree of ionisation, and the more the equilibrium is shifted to the right.
H+
H O3 +
HA + (aq) H O 2 (l) H O + 3 + (aq) A − (aq)
HA
A−
H+
Ka
+ −
(aq) HA (aq)
The equilibrium constant for the ionisation of a base is called the base dissociation constant, .
The higher the value of , the more the equilibrium is shifted right, which means that more of the base ionised, which describes a stronger base. Thus, a higher means a stronger base.
Dissociation Constant Of Water
The dissociation constant or ionisation constant of water, , has a relationship with the dissociation of an acid and the dissociation of an base.
Since and multiply to equal a constant at given temperature, and are inversely proportionate. Hence, a higher necessitates a lower
and a higher necessitates a lower .
Solution Equilibria

describe and analyse the processes involved in the dissolution of ionic compounds in water
Structure Of Ionic Compounds
Ionic compounds have a crystalline structure in which positive and negative ions are arranged in a lattice with every cation surrounded by anions and every anions surrounded by cations.
The electrostatic attraction between oppositely charged cations and anions extends throughout the whole crystal and hold the lattics together
Kb
Kb
Kb
Kw
K =w × [HA]
Ka Kb Kw Ka
Structure Of Water
Water, composed of two hydrogen atoms covalently bonded to an oxygen atom, is a polar molecule because of the large difference in electronegativity between oxygen and hydrogen atoms.
Oxygen becomes slightly negatively charged and hydrogen becomes slightly positively charged. Water is also an excellent solvent over a temperature range from 0°C to 100°C.
Dissolution Of An Ionic Compound In Water
When a solid ionic lattice is initially placed in the water, then ions begin to dissolve and forwards reaction predominates:
The dissolution process occurs due to the ion dipole interaction between ions and water. Since water is polar, the slightly negative oxygen is attracted to cation
Salt (s) Cation + (aq) Anion (aq)
and the slightly positive hydrogens attracted to anions.

investigate the use of solubility equilibria by Aboriginal and Torres Strait Islander Peoples when removing toxicity from foods, for example: – toxins in cycad fruit
Detoxifying Bush Foods
Many plant based foods the indigenous Australians ate contains toxins – substances that may cause harmful affects or even death, forcing them to demonstrate detoxification of bush foods. The first step was to cook the plant, then pounded, sliced or ground to increase surface area then leaching.
Leaching involves using water to dissolve toxins and remove it from the food. The food was placed in a flowing creek so toxins would dissolve and flow away with the river. The flowing creek serves to constantly decrease the concentration of aqueous toxin near bush food by replacing it with fresh water.
By Le Chatelier's Principle, decreasing concentration of aqeeous toxins in water would cause the equilibrium to shiift right, which turns more of the solid toxin in the food into dissolved toxin in water. Hence, dynamic equilibrium is never reached and toxins continously exit the food.

conduct an investigation to determine solubility rules, and predict and analyse the composition of substances when two ionic solutions are mixed, for example: – potassium chloride and silver nitrate – potassium iodide and lead nitrate – sodium sulfate and barium nitrate
Precipitation Reaction
A precipitation reaction occurs when two soluble substances are mixed toegther to form an insolube salt called precipitate and a soluble substance. Ions that are present but not involved in the reaction are called spectator ions. These ions make up the soluble substance at the end.
Example: when sodium chloride is added to silver nitrate
In this reaction, the precipitate formed is silver chloride while the spectator ions are sodium and nitrate.
Solubility
Solubility referse to the maximum amount of solute that can be dissolved in a given quantity of a solvent at certain temperature. To describe the solubility of
NaCl + (aq) AgNO → 3 (aq) AgCl + (s) NaNO 3 (aq)
ionic substances:
When a substance is soluble, more than 0.1 mol will dissolve in 1L of water
When a substance is insoluble, less than 0.01 mol will dissolve in 1L of water
When a substance is slightly soluble, about 0.01 - 0.1 mol will dissolve in 1L of water
Solubility Rules
To determine whether two ionic solutions will form a precipitate, the solubility rules table can be used to determine if one of the two combination of the two solutions is insolube.
Potassium Chloride & Silver Nitrate
The chemical equation for the potassium chloride & silver nitrate reaction is:
For this chemical system, the precipitate formed is silver chloride and the spectator ions are potassium and nitrate. The net ionic equation for this reaction is:
Potassium Iodide & Lead Nitrate
The chemical equation for the potassium iodide & lead nitrate reaction is:
KCl + (aq) AgNO → 3 (aq) AgCl + (s) KNO 3 (aq)
Ag + + (aq) Cl → −
(aq) AgCl (s)
2KI + (aq) Pb(NO ) → 3 2 (aq) PbI + 2 (s) 2KNO 3 (aq)
For this chemical system, the precipitate formed is lead iodide and the spectator ions are potassium and nitrate. The net ionic equation for this reaction is:
Sodium Sulfate & Barium Nitrate
The chemical equation for the sodium sulfate & barium nitrate reaction is:

derive equilibrium expressions for saturated solutions in terms of Ksp and calculate the solubility of an ionic substance from its Ksp value
Saturated Solutions
Saturated solutions is a solution that contains the maximum amount of solute that can be dissolved at a particular temperature.
Dissociated, hydrated ions are constantly precipitating into the solid substance at the same rate as the solid if going into solutions and becoming dissociated ions.
Unsaturated Solutions
Unsaturates solutions is a solutions that contains less the maximum amount of solute that can be dissolved at a particular temperature.
Insoluble Ionic Compounds
Insoluble ionic compounds are compounds that do not dissolve in water. However, it is more accurate to call these insoluble compounds sparingly soluble because only very small amounts of sparingly soluble salts dissolve and readily form saturated solutions.
The Solubility Product
The solubility product, Ksp, is the product of the concentration of ions in a saturated solution of a sparingly soluble salt. The solubility product is equal to
Pb + 2+ (aq) 2I → −
(aq) PbI 2 (s)
Na SO + 2 4 (aq) Ba(NO ) → 4 2 (aq) BaSO + 4 (s) 2NaNO 3 (aq)
Ba + 2+ (aq) SO → 4
2− (aq) BaSO 4 (s)
the equilibrium expression for a sparingly soluble salt, except only the aqueous ions are included.
For the system:
Molar Solubility
The molar solubility of a salt is the number of moles that will dissolve in water to form 1L of saturated solution.

predict the formation of a precipitate given the standard reference values for Ksp
The Ionic Product
The ionic product, , the product of the concentrations of the individual ions of a sparingly soluble salt present in a particular solution, whether it is saturated or not.
For the system:
Its ionic product if given by:
If < , then the forward reaction is favoured to reach equilibrium so the solution remains unsaturated and a precipitate will not form.
A B b a (s) aA + + (aq) bB − (aq)
K =sp [A ] [B ]+ eq a −
eq b
A B b a (s) aA + a+ (aq) bB b−
(aq)
Qsp Ksp
If > , then the reverse reaction is favoured to reach equilibrium so the solution becomes saturated and a precipitate will form.
If = , then the solution is at equilibrium so the solution becomes saturated and a precipitate will form.
The Common Ion Effect
When a salt is dissolved into a solution containing already dissolved ions where the salt added and the existing solution have a ion in common, it named a common ion. The common ion effect greatly decreases the solubility of ionic compounds - far less soluble in solutions with a common ion than they are in pure water.
If a sparingly soluble salt is added to a solution with a common ion, then it is assumed that concentration of the common ion comes entriely from the ions that are already present in the solution and not from the sparingly soluble salt.
Qsp Ksp
Qsp Ksp

module 5: equilibrium & acid reactions

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