HSC Chemistry Notes Mod 5 and 6 PDF

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Band 6 Chemistry Notes for New Chemistry Syllabus...

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Module 5: Equilibrium and Acid Reactions Static and Dynamic Equilibrium: What happens when chemical reactions do not go through to completion? Physical change: products do not include any new substances Chemical change: reactants produce new substances with different physical and chemical properties from those of the original substances

(ACSCH089, ACSCH091) Nonequilibrium systems:  

Reaction will favour movement towards greater randomness (positive entropy) Reaction will be favoured if it releases heat energy (negative enthalpy)

Equilibrium: reaction in which the rate of the forward reaction is equal to the rate of reverse reaction  

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Typically reversible o Reactants form products that react together to give the reactants back In an equilibrium, the reactants are converting to the products at the same rate that the products are converting back into the reactants o Concentration of all reactants and products remain constant over time Macroscopic properties (colour, temperature) do not change but continual microscopic change occurs between reactants and products Equilibrium can only be reached in a closed system o Closed system is a system that only exchanges energy with its surroundings but not matter and reaction contained within a certain space o Open system is system that exchanges both energy and matter with its surroundings and substances can be added or lost

Static Equilibrium:  

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Reaction occurs in one direction and reactants continue to produce the products until one of the reactants is used up and the reaction stops Rate of forward reaction and rate of reverse reaction is zero o No exchange between reactants and products and reaction o Irreversible Eg diamond  graphite o Exists in equilibrium but activation energy for diamond to turn back into graphite is so high that rate of reverse reaction to be taken is zero

Dynamic Equilibrium:

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Rate of forward reaction and the rate of reverse reaction are equal but not zero o Exchange between reactants and products and reaction o Reversible Eg hydrogen and nitrogen to ammonia o 3H2 + N2 ⇌ 2NH3

Enthalpy (H): total amount of heat in a system with constant volume and pressure

∆ H=mC ∆ t H = change in enthalpy or heat absorbed or released (Joules, J) C = specific heat capacity t = change in temperature (Celsius, C or Kelvin, K) m = mass of substance (g)  

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Changes in enthalpy from a chemical reaction can be calculated ∆ H > 0 ENDOTHERMIC REACTION o Positive change in enthalpy o Heat absorbed ∆ H < 0 EXOTHERMIC REACTION o Negative change in enthalpy o Heat is released

Entropy (S): measure of disorder or randomness within a system  

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The tendance of the molecules to spread out due to higher probabilities of different configurations (or microstates) a system can be in Entropy increases with o Higher temperature  More kinetic energy  More ways to spread out o Greater number of particles  More particles to form different combinations o Greater complexity of compounds  More variance o Greater particle movement  More freedom of movement with more space Change in entropy can be determined from analysing the balanced chemical equation of reaction o ∆ S > 0 INCREASE in entropy ∆ S < 0 DECREASE in entropy o

Gibbs Free Energy

∆ G=∆ H °−T ∆ S ° G = change in Gibbs Free Energy (kJ/mol) T = temperature (Kelvin, K)  

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Quantity that analyses effect of drivers to determine whether reaction will occur spontaneously at a given temperature Combined effect of enthalpy and entropy changes represented at constant temperature and pressure o Calculates total amount of energy available in a system If free energy decreases, reaction is proceeds however, if the free energy increases the reaction cannot proceed o ∆ G < 0 EXERGONIC  Reaction is favoured and spontaneous  Loss of energy with the reaction proceeding ∆ G > 0 ENDERGONIC o  Reaction not favoured and nonspontaneous  Requires the input of energy to occur  Reverse reaction would occur spontaneously For nonequilibrium systems, Gibbs free energy is negative for one side of the reaction and positive for the other side of the reaction o Typically a system wants to increase S and decrease H (and going against this requires energy) o With ∆ H < 0 + ∆ S > 0 = ∆ G < 0 Eg combustion o Exothermic  Decrease in enthalpy  Increase in entropy o Always spontaneous, whilst reverse reaction is nonspontaneous and is a nonequilibrium reaction Eg photosynthesis o Endothermic  Increase in enthalpy  Decrease in entropy o Nonspontaneous, forward reaction (respiration) is spontaneous and is a nonequilibrium system

(ACSCH070, ACSCH094) Collision Theory and Equilibrium Reactions 

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Explains and predicts the rate of chemical reactions o Atoms and molecules must collide in order to react o The RoR depends on number of effective collisions o An effective collision depends on energy and orientation of collision Effect collisions results in a chemical reaction through the breaking and reforming of chemical bonds

Sufficient energy to break chemical bonds  Kinetic energy of molecules o Correct orientation for effective collision Activation energy: minimum energy required to break the bonds to initiate a chemical reaction o Once energy available is equal or greater than activation energy, the correct orientation of particles is required for to break the bonds Orientation o Effective collision at break the bond, rather than bouncing off molecules Once equilibrium is established, the RoR can be affected o Temperature  Increases kinetic energy for more effective collisions due to increase in molecules with sufficient activation energy o Pressure / Concentration  Increases density of molecules to increase collision frequency o SA  Increase SA increases collision frequency o Presence of catalyst  Provides alternative pathway with lower activation energy  Allows for more effective collisions o

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Factors that Affect Equilibrium: What factors affect equilibrium and how? Le Chatelier’s Principle 

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Le Chatelier’s Principle states if a system at equilibrium is disturbed, the system adjusts itself to minimise the disturbance o Effect of disturbance is not completely removed but minimised Equilibrium position affected by o Concentration o Temperature o Pressure and volume  Inverse relationship (P increases with decrease in V) Concentration o A+B ⇌ C+D o Increase of compound A will lead to shift in system to minimise increase, favouring the forward reaction to convert compound into more of C + D o At the new equilibrium, concentrations remain consistent but not the same as prior to disturbance o On a graph  Sharp increase (addition of compound) followed by gradual decrease  Gradual increase of other Temperature o The effect of change depends on the nature of the reaction as endothermic or exothermic o Increase in temperature would favour the endothermic reaction as heat would be absorbed to minimise increase in heat

Decrease in temperature would favour exothermic reaction as heat will be released to minimise decrease in heat o On a graph  Gradual increase and decrease Pressure o Pressure only affects gases, as liquids and solids cannot be compressed o Effect depends on the number of moles of gas in the equation  Increase in pressure will result in results in shift towards sides with fewer moles of gas  This is due to the nature of one mole of gas taking up the same volume as one mole of any other gas, and by reducing the number of moles the pressure is reduced o If there is an equal number of moles of gas on both sides of equation, and the volume/pressure is changed, then concentration of both change equally and equilibrium has not been altered o On a graph  Drastic drop in both  Continued gradual decrease in one  Gradual increase in other Volume o Any change in volume results in a change in pressure o Volume of system determines number of moles of gas in space o Increase in volume decreases pressure o Decrease in volume increases pressure Catalyst o Substance that increases the rate of reaction without being used up in the reaction o Lower activation energy by providing alternative pathway o Does not have effect on equilibrium  Rates of both reactions increase, but change is by same proportion o

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Activation Energy and Heat of Reaction on Equilibrium  

Activation energy and heat of reaction (enthalpy change) affect the position of the equilibrium Represented by energy profile diagrams

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Enthalpy (heat of reaction) is the same but opposite in heat transfer

Calculating Equilibrium Constant: How can the position of equilibrium be described and what does the equilibrium constant represent? (ACSCH079, ACSCH096) Equilibrium Constant

B ¿ ¿ [ A ]a ¿ c D [ C] [ D ] K eq = ¿  

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Measures the proportion of equilibrium based on ratio of concentration of reactants and products [ ] show the concentrations of substance o C + D as product o A + B as reactants Indices as the coefficients in front of each substance in equation Only species that are gases or solutes appear in equilibrium expression o Pure solids and liquids do NOT appear o Eg water Constant K is a value that indicates the position of the system at equilibrium o How far to completion it is The larger the constant, the more the equilibrium favours the products (forward reaction) o K > 1, products favoured  K > 103, equilibrium almost proceeds to completion o K = 1, significant concentrations of both reactants and products present at equilibrium o K < 1, reactants favoured

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K < 10-3 equilibrium almost doesn’t proceed to completion

Effect of Temperature on K eq 

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Temperature is the only factor that changes the value of the equilibrium constant o With concentration, volume or pressure changes in equilibrium, the concentrations of reactants and products adjust so that the numerator and denominator retain same ratio prior to change Exothermic reactions o Lowered temperatures increase K  Reaction favours the exothermic reaction to increase heat  Creates more product, increasing the numerator o Increased temperature decrease K Endothermic reactions o Increased temperature increases K  Reaction favours endothermic reaction to decrease heat  Creates more product, increasing numerator o Decreased temperature decreases K Equilibrium constant not fixed but dependent on the system

Reaction Quotient 

When a reaction is not at equilibrium or unsure of equilibrium, Q replaces K in the same equation

B ¿ ¿ [ A] a ¿ c D [C ] [D ] Q= ¿  

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Indicates relative amount of product and reactant in reaction at given time Q K o More products and less reactants than at equilibrium o System lies too much to the right (products) o Reverse reaction favoured

Dissolution of Ionic Compounds 

Ionic compounds are a group of atoms that are chemically bonded together by electrostatic attraction between oppositely charged ions

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At standard room conditions, ionic compounds exist in solid crystalline lattice structure Most ionic compounds are soluble in water o When introduced to liquid water, crystalline structure breaks down as ions dissociated in solution o When soluble ionic compound added to water, water molecules surround the one ion within lattice due to electrostatic attraction between polar water molecule and the charged ion o Negative end (oxygen) is attracted to cations whilst positive end (hydrogen is attracted to anions), forming ion dipole interactions When ion dipole force is greater than electrostatic force between ions in lattice, ions are dislodged from position in crystal o Completely removed from lattice to be surrounded by water molecules and become hydrated or solvated Salvation layer acts shield and prevents solvated anion from colliding directly with solvated cation and keeps solvated ions in solution

(ACSCH098, ACSCH099) Dissociation of Acids   

In weak acids (partial ionisation), some of their molecules break apart to form hydrogen ions and anions of the acid Other molecules of the acid remain as molecules and do not ionise, forming an equilibrium Since concentration of water is approximately constant in a dilute solution, the equilibrium constant for the ionisation of an acid is called the acid dissociation constant (K a)

+¿ ¿ H3O ¿ −¿ ¿ A ¿ ¿ K a=¿

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The size of constant provides information about the degree of ionisation that occurs o The larger the Ka the larger the value of the numerator in the general equation for K eq and hence, greater the concentration of ions compared to molecules of acid o Basically, the larger the Ka the greater the degree of ionisation

Dissociation of Bases 

Weak base is when only some of the molecules of base react with water to produce ions

+¿ ¿ BH ¿ −¿ OH ¿ ¿ ¿ K b=¿

Equilibrium Constant for Gaseous Systems in Pressure 

Mole fraction: determine for each species when there is a mixture of gases

mole fraction= 

number of moles of species total number of moles of gas present

Partial pressure: proportion of pressure that is due to collisions for particular gas present

∂ pressure=mole fraction of species × total pressure of system

a

K p=   

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P A × PB c d PC× PD Brackets not used as they indicate concentration Only gaseous species used Equilibrium constant for pressure is different to equilibrium constant for concentration

Solubility Equilibria: How does solubility relate to chemical equilibrium? 

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Structure of ionic compounds o Crystalline structure o Positive and negative ions arranged with every positive ion surrounded by negative ions and vice versa o Electrostatic attraction between oppositely charged ions o Cations and anions held together by ionic bonds Structure of water o 2 hydrogen atoms covalently bonded to oxygen atom o Polar molecule with large difference in electronegativity between O and H  O becomes slightly negatively charged as strong electronegativity attracts electrons in OH covalent bond are more attracted to oxygen  H is slightly positively charged o Water displays hydrogen bonding o Water acts as a solvent by forming hydrogen or ion-dipole bonds with substances Water of crystallisation o Occurs when water molecules are attracted to ions of salt o Bonds between ions and water are ion-dipole bonds  Dipole in water molecules form weak bond o Water or molecules that form dipole bonds with metal atom are ligands

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 Ligands take positions around central ion following VESPR theory When all water molecules have been removed, compound is anhydrous

Dissolution of Ionic Compounds 

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When salt dissolves, ion-dipole bonds are formed o Positive H of water attracts anion of salt o Water molecules surround anion until it is hydrated o Negative part of water is attracted to cation and undergoes similar process until ionic bonds within salt crystal and broke and the salt has dissolved For ionic substances to dissolve, the energy required to separate ions has to be less than energy released when ions are hydrated o Energy required to separate depends on strength of ionic bonding  This depends on the arrangement of the ions in lattice and charges of ion o Energy released with hydration depends on strength of ion-dipole attraction  This depends on size of ion, charges of ion and geometry of ions (polyatomic) o The more energy released compared with energy needed, the more soluble the salt

Solubility: maximum mass that can dissolve in 100g of solvent at given temperature Unsaturated solution: solution where more of solute can be added and dissolved Saturated solution: point where no more solute can be added Supersaturated solution: more solute dissolved than in a saturated solution at the same temperature

Solubility Product Constant 

Solubility of ionic compound is determined by amount it dissolves in a solution before becoming saturated o Solution is saturated when it contains maximum concentration of solute at given temperature o Excess solute is precipitated out as a solid o Saturated solutions exist in dynamic equilibrium, where rate of dissolution equals rate of precipitation c

K sp=[C] [ D ] 

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D

The reactant is not included in the equation as it is a solid (dissolution reaction), and it does not have measurable concentration o Heterogenous system Ksp represents the level at which a solute dissolves in solution o Higher Ksp, the more soluble a substance is Equilibrium systems only occur for ‘sparingly soluble’ or ‘insoluble’ salts

Formation of Precipitate

Q sp =[C ]c [ D ] D 

Used to predict whether a precipitate will form using ion product constant Q o Q < K, precipitate will not form (unsaturated) o Q = K, precipitation will not form (saturated) o Q > K, precipitation will form (supersaturated)

Solubility Equilibria by Aboriginal and Torres Strait Islanders 

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Cycad fruits as important food source, however, is highly toxic and carcinogenic o Contains poison cycasin o Induced vomiting, diarrhea, weakness and seizures Leaching  for soluble toxins o Cutting open kernels and immersing into water o Toxins leach out into water and dissolve it from plant in basket placed in a flowing creek so toxin would flow away with water o Cycasin as soluble in water to allow toxins to diffuse and be washed out Cycads can cause gastrointestinal and liver disorders, including cancer and also strip insulating layer of myelin from around the nerve, causing nerve impulses to leak for unpredictable movements or even death

Solubility Rules  

Precipitation reaction: solutions of certain ionic compounds reacting to produce a solid or precipitate Spectator ions: ions not involved in reaction

Nitrates

Exceptions:

Acetates

* PMS

^ Castro Bear

Group 1

Pb

Calcium

Sulfates *^

Mercury

Strontium

Ammonium

Silver

Barium

Group 17 *

Common Ion Effect  

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Occurs when another chemical has ion in common with original ionic salt added to solution Increased concentration of shared ion will lead to LCP change o Chemical system adjusts to overcome increase in concentration of shared ion o Will favour reverse reaction (solid) over the dissociated ion (shared) Thus, the salt will be more insoluble with common ion

Module 6: Acid Base Reactions Properties of Acids and Bases: What is an acid and what is a base? (ACSCH067) IUPAC of Inorganic Acids H+ bonded to Non-Oxygen Element (Binary Acid)    

Name of acid begins with prefix hydroEnds with suffix -ic Root of name of anion in between Examples o Hydrofluoric acid o Hydrochloric acid o Hydrobromic acid o Hydro sulfuric acid

H+ Bonded to Oxygen Atom of Polyatomic Anion (Oxoacid)  

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If anion ends in -ate, then name of acid replaces suffix with -ic o Slight variation for sulfuric and phosphoric acid If anion ends in -ite, then name of acid replaces suffix with -ous o Has less oxygen o Eg sulphate  sulf...