hsc head start lecture chemistry · 2020. 4. 25. · lecture overview i’ll be going through all...

HSC Head Start LectureCHEMISTRY

Presented by:JIANNE PARISI

• Session 1– Who am I?– Our goals today– Static and dynamic equilibrium

• Session 2– Factors that affect equilibrium– Calculating the equilibrium constant

• Session 3– Solution equilibria– General HSC tips

Lecture OverviewI’ll be going through all of module 5!

• I graduated high school in 2018 – I studied advanced English, extension maths, chemistry, biology, music 2 and music extension and received 2 Encore nominations!

• I’m now studying a double degree of commerce and advanced computing at Usyd

• My HSC experience was quite awful ngl!• But I would like to help make yours less bad

Who am I?

• Get a good basic understanding of module 5• Feel confident answering questions• Learn good study strategies• Take lots of notes• Ask questions!!!

Our Goals Today

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

• So far you have only seen irreversible reactions• These go to completion and use the normal (à) arrow• This means that reactants become products and that’s all!

• This module familiarizes you with reversible reactions• Reactants become products but products can also become reactants• Reversible reactions use this arrow: ⇌ (it makes sense)

• For reversible reactions, changing conditions will affect the proportion of reactants to products.

5.1 Static and Dynamic Equilibrium

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Practical investigations (reversible)Cobalt(II) chloride hydrated and dehydrated• CoCl2(aq) + 6H2O(l) ⇌ CoCl2 · 6H2O(l)• Hydration/dehydration = adding or removing water• Cobalt(II) chloride becomes cobalt(II) chloride hexahydrate• There is a visible difference! The anhydrous (dehydrated) form is sky

blue and the hydrated form is purple

Iron(III) nitrate and potassium thiocyanate• Fe3+(aq) + SCN-(aq) ⇌ FeSCN2+(aq)• Iron ions are yellow in solution and potassium thiocyanate is clear but

iron thiocyanate is BLOOD RED


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Practical investigations (irreversible)Burning magnesium• 2Mg(s) + O2 à 2MgO(s) + energy• Note the arrow! This is irreversible• This reaction releases A LOT of energy – we can sometimes reverse

reactions by putting as much energy back into the system but this is completely impractical for combustion because it releases so much

Burning steel wool

• 2Fe(s) + !"O2(g) à Fe2O3(s)

• Works the same way as burning magnesium


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Equilibrium• A state where concentrations of reactants and products stay

approximately constant

Dynamic Equilibrium• Over time, reversible reactions will reach a state called dynamic equilibrium• Conditions for dynamic equilibrium:

– Reaction is reversible– Closed system– Rate of the forward reaction equals the rate of the reverse reaction– Concentrations of reactants and products are constant but not necessarily equal– Macroscopic properties are constant

• While concentrations are staying constant, individual particles are still oscillating between reacted and unreacted states! That’s why this equilibrium is called dynamic.

• The rate at which particles shift backward and forward is equal


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Static Equilibrium• Static because all particles are at rest• They do NOT oscillate between reacted and unreacted states• For example any irreversible reaction – reactants become products and cannot

switch back to reactants• It is still an equilibrium because concentrations of reactants and products

don’t change once the reaction has gone to completion• The rate of forwards and backwards reactions are zeroOpen vs closed systems• Whether a system can or cannot interact with its surroundings• If a system is open, the external environment will have an impact on equilibrium

being reached• E.g. mixing hot and cold water:

– If in an insulated box, the water will be warm and NOT affected by changing the temperature of the environment

– If in an open box, the water will be warm and then cool down until it reached the temperature of the environment


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Non-equilibrium SystemsEntropy• How many possible states can something be in?

– Solids – stuck in a lattice– Liquids – some freedom of movement but they’re still kinda stuck together– Gases – lots of freedom of movement – most possible states bc individual things

have the largest range of possible positions

• This can be looked at as how disordered something is• If there are more states, things are more likely to be in one of these

states• Entropy is always increasing for the universe• Just things moving to a more statistically probable arrangement, which

will tend to be more disordered (toward higher entropy)


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Non-equilibrium SystemsEnthalpy• Enthalpy – heat content of a system• Chemical energy is stored in bonds and can be released as

heat• Conversely, heat energy can be stored as chemical energy

in bonds• For an endothermic reaction, enthalpy with INCREASE

– The reaction takes energy IN• For an exothermic reaction, enthalpy will DECREASE

– The reaction sends energy OUT


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Entropy and enthalpy in non-equilibrium systemsCombustion reactions• C2H5OH(l) + 3O2(g) à 2CO2(g) + 3H2O(g)• Combustion reactions are entropically favoured because the entropy of the

system increases• There are more products than reactants so the system can exist in more

possible states (entropy goes UP)• Going from gas and liquid to just gas – gases can exist in more states• Enthalpy decreases because heat from the system is lost to the

environmentPhotosynthesis• Photosynthesis is NOT entropically favoured – entropy goes DOWN

– 6CO2(g) + 6H2O(l) à C6H12O6(aq) + 6O2(g)– Goes from more molecules to fewer molecules

• And it is endothermic so enthalpy goes UP• It requires energy from the sun in order to proceed


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Module 5: Equilibrium and Acid Reactions

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Relationship between collision theory and reaction rateCollision Theory – in order for a successful reaction between gas particles:• A collision must occur• It must have sufficient speed• Particles must be in the right orientationActivation Energy• There is a minimum energy required for a reaction to occur. Granted that

particles are in the right orientation, a collision with sufficient energy (i.e. with sufficient speed) will be successful

Reaction rateA reaction will proceed faster if there are more collisions above the activation energy.• Increasing concentration of reactants increases the likelihood of a collision

happening, and of it being in the correct orientation• Increasing the temperature of reactants increases their kinetic energy, thereby

increasing their speed, so the likelihood of a collision being above the activation energy is increased


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5.1 Practice Questions

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What factors affect equilibrium and how?• Temperature• Concentration• Volume/pressure (FOR GASES)

Le Chatelier’s Principle• “A system at equilibrium, when introduced to a change, will shift in

equilibrium to minimise that change.”• So a system will favour either the forward or reverse reaction in order to

get comfy again

So for example: If I increased the concentration of a particular substance, the system would favour the reaction that decreases its concentration.

5.2 Factors that Affect Equilibrium

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Concentration• If the concentration of the reactants is increased, the

equilibrium will favour the reaction that decreases the concentration of reactants (remember equilibrium shifts to minimise a change according to Le Chatelier’s Principle)

• The forward reaction is favoured (uses up reactants and makes more products)– This is called shifting to the right as more products are produced

• Conversely, if we increase the products, the equilibrium will shift to the left– Why?


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Notes:• Reactants are on the left and products are on the right

– You shouldn’t switch the names around for the reverse reaction! The reverse reaction involves products becoming reactants

• The system doesn’t want to get back to exactly where it was – it wants to find a new equilibrium under the new conditions and this new equilibrium will either be to the left or right of the previous equilibrium

• Shifting to the left = concentration of reactants goes up and concentration of products goes down

• Shifting to the right = concentration of products goes up and concentration of reactants goes down


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Pressure• Only affects GAS (!!!!!!!)• If we increase the pressure of a system, it will shift to

decrease pressure by favouring the side with fewer moles of gas– Each mole of gas (any gas) occupies the same volume

• If we decrease the pressure of a system, it will shift to increase pressure by shifting to the side with more moles of gas


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TemperatureRecall:• Endothermic reactions absorb heat to move forward

– It might help to think about heat as a reactant• Exothermic reactions release heat to move forward

– it might help to think about heat as a product• If a forward reaction is endothermic, the reverse reaction will be

exothermic and vice versa

• If we increase the temperature of a system, then by Le Chatelier’sPrinciple, the system will shift to minimize this change and lower the temperature of the system by favouring the endothermic reaction

• If we decrease the temperature of a system, it will shift to minimize this change and increase temperature by favouring the exothermic reaction


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In practice:• Heating cobalt(II) chloride hydrate:

– CoCl2(aq) + 6H2O(l) ⇌ CoCl2 · 6H2O(l)– We know that warming it up will cause the water to evaporate, dehydrating the

system– This means that the equilibrium shifts to the left when there is an increase in

temperature– So the reverse reaction must be endothermic (uses up heat)– And that means that the forward reaction is exothermic– We know that the dehydrated form is sky blue and the hydrated form is purple – SO heating the mixture will make the solution turn blue


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In practice:• Interaction between nitrogen dioxide and dinitrogen tetroxide

– 2NO2(g) ⇌ N2O4(g) ΔH = -57.20 kJ/mol– What happens if we increase concentration of NO2 or N2O4? (equilibrium shifts away

from the increase)– What happens if we increase the pressure of the vessel? (equilibrium shifts to the

right to decrease pressure because there are fewer moles of gas)– ΔH is negative so the reaction is exothermic. What would happen if we increased

temperature (equilibrium shifts to the left to decrease temperature)– To maximise yield: continuously add more reactant, continuously remove product,

increase pressure, decrease temperature

• iron(III) thiocyanate and varying concentration of ions– Fe3+(aq) + SCN-(aq) ⇌ FeSCN2+(aq)– Adding silver nitrate will cause a silver thiocyanate precipitate to form, reducing the

concentration of SCN- in solution and causing the equilibrium to shift to the left to increase concentration. Solution should become less red.


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Observations about equilibrium in terms of collision theory• Higher concentration = more reactant particles = more collisions• Higher pressure basically increases the concentration of gases• Higher temperature = more kinetic energy = more collisions that are

above the activation energy

How do activation energy and heat of reaction affect the position of equilibrium?• Lower activation energy means it’s easier for a reaction to take place• The higher the activation energy, the more kinetic energy is required in

order for a reaction to take place• Endothermic reactions are favoured when temperature is increased and

exothermic reactions are favoured when temperature is decreased


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

aA +bB ⇌ cC + dDKeq =

C - D /A 0 B 1

• Square brackets mean concentration (specifically use concentration at equilibrium)

• Turn coefficients into exponents and the multiple everything

• High K = lots of product• Low K = not as much product

Homogenous reactions?• A reaction that occurs in a single phase – all reactants and products are of

the same state

5.3 Calculating the Equilibrium Constant

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Q vs KYou get Q when you use the K calculation but the system isn’t at equilibrium – Q means reaction quotient

• Q < K = not enough product = favours the forward reaction to reach equilibrium

• Q > K = too much product = favours the reverse reaction to reach equilibrium


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Equilibrium math: using ICE tablesFinding other concentrations in a system given one concentration• Write out the equation to get molar ratios• 2HI(g) ⇌ H2(g) + I2(g) (this is 2:1:1)• So 2 moles of HI will produce 1 mole of H2 and 1 mole of I2

• Hypothetically, if we let 2 moles of HI decompose in a 1L vessel, waited a while and then found that there were 0.7 moles of I2 at equilibrium:– There would also be 0.7 moles of H2 because H2:I2 is 1:1– An increase of 0.7 moles of I2 or H2 equals a decrease of twice as much HI

(remember it takes 2 moles of HI to produce 1 mole of either product)


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HI H2 I2Initial 2 moles 0 0

Change -1.4 +0.7 +0.7

Equilibrium 0.6 0.7 0.7

Temperature• The only factor that can change K• Remember:

– Increasing temperature in an exothermic reaction will decrease Keq– Increasing temperature in an endothermic reaction will increase Keq

Uses of Keq

• Keq = %&'()(*+ (&, [,+./)(*+ (&,][(&,(1 1&2%&3,4]– Dissociation of ionic compounds– This will tell us how much something dissolves in water

• Ka = 67 [89][86]

– Used to describe how the degree to which an acid ionises


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How does solubility relate to chemical equilibrium?

Dissolution of ionic compounds in water• Oxygen has 2 lone pairs in water molecules• O is very electronegative (hogs the electrons in the covalent bonds),

leaving H quite positive• So water is a polar molecule• Positive ions are attracted to O and negative ions are attracted to H,

forming ion-dipole bonds

5.4 Solution Equilibria

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Removing toxins in cycad fruit• Cycad fruit is highly toxic and carcinogenic• Indigenous peoples needed a method for reducing the toxicity• One of these methods was to cut open a fruit and leave it in running

water• The toxins would dissolve into the water: Toxin(s) ⇌ Toxin(aq)• Moving water means that the aqueous toxin is continually carried away

and, continually decreasing the concentration of aqueous toxin• SO equilibrium is constantly shifting to the right and the toxins in the

fruit decrease


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Solubility Rules• A useful mnemonic (credit to Jacob Silove)

• More active metals will displace less active metals from solution• The table of standard potentials on your reference sheet is in order from

most to least active– E.g. KCL(s) + AgNO3(aq) ⇌ AgCl(s) + KNO3(aq)


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Mixing ionic solutions• Potassium chloride and silver nitrate

– KCl(aq) + AgNO3(aq) ⇌ AgCl(s) + KNO3(aq)– Silver chloride is insoluble (it’s a white precipitate)

• Potassium iodide and lead nitrate– Pb(NO3)2(aq) + 2KI(aq) ⇌ PbI2(s) + 2KNO3(aq)– Lead iodide is insoluble (it’s a yellow precipitate)

• Sodium sulfate and barium nitrate– Ba(NO3)2(aq) + Na2SO4(aq) ⇌ BaSO4(s) + 2NaNO3(aq)– Barium sulfate is insoluble (it’s a white precipitate)


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Ksp (solubility product)• Ksp = [A-][B+]• No denominator because it’s a solid!• Saturated solution

To get from Ksp to solubility• Solubility = grams dissolved/volume of water dissolved into• Ksp is the product of concentrations of the products

– Use it to work out moles of products

• And then use moles of product to work out moles of reactant (solubility is based on the initial substance you were trying to dissolve)

• Convert moles of reactant to grams of reactant using m = nM• Solubility = grams per litre


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Predicting if a Precipitate will formA precipitate forms when the quantity of a soluble substance exceeds a saturation point

• Find the limiting reagent• Calculate the solubility (from Ksp) (this is the saturation

point)• Determine whether you have tried to dissolve more than

can physically dissolve (more grams than the solubility allows)


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CaSO4(s) ⇌ Ca2+(aq) + SO42-(aq)

H2SO4(aq) ⇌ 2H+(aq) + SO42-(aq)

Because more sulfate ions are added to the equilibrium, by Le Chatelier’s principle, the system will shift in the direction that decreases concentration of these ions. (ignore the hydrogenions bc they don’t matter for this question) So the reverse reaction is favoured (equilibrium shifts to the left), calcium and sulfate ions get used up and converted into solid calcium sulfate, meaning that there are fewer ions in solution


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Solution


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Why chlorine ions have a charge of -1


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Cheeky shortcut to make calculation easier: just get rid of one x to isolate the other

(given)

(from your table)

The x has just been disregarded (it’s so minute it doesn’t make enough of a difference to matter)


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(molecular mass of calcium hydroxide is 74.093)


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MODULE 5 DONE

• Scientific drawings– Don’t close things that aren’t closed– Don’t shade

• Graphing– As always, independent variable on the x axis– Use at least 50% of the space

• Mole calculations– If you get lost, just write down each of the values you’ve been given

and then all the calculations you actually can do (even if you don’t get anywhere, this is a good way to scavenge marks)

Skills

GET A GOOD STARTBecause catching up is 1000000000 times harder than just

being on top of things from the beginning

You will have 4 internal assessments that make up 50% of your overall mark

It ultimately doesn’t matter if you fuck one up!

DON’T BE LAZYBurnout is real but so is laziness. Find a balance and be

honest with yourself. It’s ok if you can’t do more. It’s not ok if you can but you can’t be bothered

HSC Tips

WRITE NOTES AS YOU GONo, you canNOT do it later there will not be time

I’d advise that you do them electronically so you can go back and edit! And then print them, highlight them, draw on them

and get to know them like a really good friend

AVOID CRAMMING WHEN YOU CANBy doing weekly recaps! Each weekend, go over what you

learnt that week and also what you learnt the previous week and it will make exam revision a shitload easier

HSC Tips

FOLLOW THE SYLLABUSThings are much less overwhelming if you know wtf is going on. Also the syllabus is there for a reason! Stuff that isn’t in

the syllabus isn’t assessable

PLAN YOUR TIMEWrite things down for goodness sake or you will forget to do

them. Keep to-do lists and make sure you PRIORITISE

PRIORITISEI can almost guarantee you won’t have time for everything

unless you’re a literal robot. Things that are assessable always come first (but not before your mental health)

HSC Tips

CHANGE UP YOUR STUDY ENVIRONMENTSometimes it will be really hard to focus! Try studying outside,

studying lying on the ground, studying at other people’s houses, studying on the toilet to give your brain the change it wants

TAKE BREAKSDon’t waste time bashing away at something when you can’t

focus! Go for a walk or do anything but study for a set amount of time to reset and you’ll ultimately save time

USE A STUDY/BREAK SPLITYou can go as small as 5 minutes of study to 10 minutes of

break, it’s better than getting nothing done at all!

HSC Tips

THE INTERNET IS FULL OF USEFUL STUFFYou can learn almost anything by googling it. You’ll only get

stuck when you can’t be bothered to do it

DON’T MURDER YOUR SOCIAL LIFEBecause studying is going to drive you crazy and your friends

will keep you sane

STUDYING = SOCIALISING?Study with your friends! Not only do you get to spend time with

them but teaching is the best way to learn and many heads reduce the workload

HSC Tips

WORK ACCORDING TO YOUR GOALSThere is no point in working your butt off and stressing if you

don’t even need a high ATAR, especially because after the HSC, you’re never going to use a good 90% of what you learnt. Get the ATAR you need and move on (no one cares about your

ATAR anyway)

YOU GET OUT WHAT YOU PUT INYou can’t do minimal work and expect a high ATAR. The HSC generally favours people who work hard over people who are

naturally smart.

HSC Tips

IT’S ONLY ONE YEARYes it’s shit but you only get so much time to do as well as you

can and it has a definite end (wahoo!!!)

DON’T BE ANGRY AT YOURSELF FOR THINGS YOU HAVE NO CONTROL OVER

This is probably at the top of the list of completely useless and quite destructive things

TAKE CARE OF YOURSELF AND DON’T GIVE UP!The HSC is HARD but it doesn’t have to be the most stressful

experience of your life

HSC Tips

• Everyone is in the same boat. At the time you’re crying there is almost definitely someone else also crying.

• Sometimes it’s helpful to find someone to cry with/at and you’d be surprised by how many people are willing.

• The best way to minimize crying is to do things early BUT you can’t change the past and being angry at yourself solves nothing so get over it and make the most of the time you’ve got.

• The best you can do is the best you can do.• Your ATAR is an easy way into a course but it’s a perfectly valid plan to

do a similar course if you don’t get into your first choice, take the units you’d be taking for the course you want and have them count as credit when you transfer so you don’t even lose that much time (not that losing time should even matter)

• (You’ve heard it a million times but) Other than making it into a course, your ATAR counts for nothing and no one gives any fucks about it!

Emotional Support Slide

hsc head start lecture chemistry · 2020. 4. 25. · lecture overview i’ll be going through all...

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