CHEM Revision!!

xD
Legend
● e- means electron
● ox.no. means oxidation number
● En means Electronegative (I think?)
● EC series means Electrochemical series
Mole Concept
(exclude the empirical & molecular formula; what are they lol i forgot)
(ooh empirical is the one where you draw a table and find the ratio of the
elements in the compound or sth so like C2H3O200000 or something)
Molecular formula is the final compound C4H6O400000 like empirical*number

Relative atomic mass (Ar)

- Average mass of atom in relation to 1/12
mass carbon-12 atom
Relative molecular mass (Mr)
- Basically add up mass of all the atoms in the
molecule
Mole Concept - Percentage thing
(Guys i changed it fyi)
Percentage purity

Expected yield (pure)

------------------ ----------- x 100%
Actual yield (impure)
 Mole Concept-
percentage mass composition

Ar of element x no. of atoms

--------------------------------------- x
100%
Mr of compound % mass of H in H2O2 = (2.0 ÷ 34.0) x 100%
Mr of H2O2 = 2(Ar of H) + 2(Ar
= 5.9%
of O)
% mass of O in H2O2 = (32.0 ÷ 34.0) x100%
= 2(1.0) + 2(16.0)
= 34.0 = 94.1%
 Mole Concept
Avogadro’s number, NA = 6.02 x 1023/mol
1 mol of anything contains 6.02 x 1023 particles of that
substance

Molar mass (g/mol)

Molar Gas Volume (gas volume per mol thing):

- R.t.p., room pressure (1 atm): 24dm3/mol or 24000cm3/mol
- S.t.p. (0ºC), room pressure (1 atm): 22.4dm3/mol
 Mol

(Copied from last year’s slides lol)

Mol
Concentration
● g/dm3
● mol/dm3

mol ratio = gas volume ratio

Stoichiometry
Hey what is this
satanic ritual
Bey that i am not
The ritual of rhyme O_O that will grant and
secure and ensure you a GPA of 5.0 for all of
part of eternity

Day
Hay
May
Stoichiometry
Just calculate the amount (no. of mol) and
mass of certain compound in an equation

● g/mol x mol = g
● g / g/mol = mol
● Gas volume stuff
QA (copied from QA notes lel sorry
for tiny font)
Anion tests
Anion Test Test Results
Carbonate (CO32- Add dilute acid Effervescence observed, CO2
) produced
Chloride (Cl-) Acidify with dilute acid, then add aq White ppt
AgNO3
Iodide (I )
- Acidify with dilute HNO3, then add aq Yellow ppt
Pb(NO3)2
Sulfate (SO42-) Acidify with dilute aq HNO3, then add White ppt
Ba(NO3)2
Nitrate (NO3-) Add aq NaOH, then aluminium NH3 produced
foil/powder or Devarda’s alloy; warm
carefully
Solubility Table
Soluble Insoluble
Salts of NH4+, K+, Carbonates except Na+, K+, NH4+
Na+
Nitrates Hydroxides Slightly soluble:
Hydrogen Oxides CaO, Ca(OH)2, CaS, CaSO3
carbonates
Sulfates except PbSO4, Sulfides (S2-)
BaSO4, CaSO4
(slightly soluble)
Halides including: Sulfites (SO32-) ^personally i think it’s
kinda useful in
chlorides except PbCl2, AgCl
understanding the
bromides except PbBr2,
cation tests
iodides AgBr
except PbI2, AgI
 QA Cation tests
Cation Effect of aq NaOH Effect of aq NH3
Aluminium (Al3+) White ppt, soluble in excess giving a colourless solution White ppt; insoluble in excess (Add Halide/Sulfate)
Add HCl (aq)
Lead(II) (Pb2+) White ppt, soluble in excess giving a colourless solution White ppt; insoluble in excess White ppt: Pb2+ (PbCl)
Colourless: Al3+ (AlCl3)

Add KI (aq)
Yellow T00ppt: Pb2+ (PbI2)
White ppt: Al3+

Add NaSO4 (aq):

White ppt: Pb2+ (PbSO4)
Colourless: Al3+ (Al2(SO4)3)
Zinc (Zn2+) White ppt, soluble in excess giving a colourless solution White ppt; soluble in excess giving a colourless solution

Calcium (Ca2+) White ppt, insoluble on warming No ppt

Ammonium (NH4+) Ammonia produced on warming -
Copper(II) (Cu2+) Light blue ppt, insoluble in excess Light blue ppt; soluble in excess giving a dark blue solution
Iron(II) (Fe2+) Green ppt, insoluble in excess Green ppt, insoluble in excess
Iron(III) (Fe3+) Red-brown ppt, insoluble in excess Red-brown ppt, insoluble in excess

NH4+ No visible reaction Test for ammonia -

Na+ No visible reaction No visible reaction (flame test not tested)
K+ No visible reaction No visible reaction
(flame test not tested)
Ca2+ White ppt, insoluble on warming No visible reaction NaOH test for Ca
QA
Test for gases

Ammonia (NH3) Turns damp red litmus paper blue

Carbon dioxide Gives white ppt with limewater, ppt dissolves with excess
(CO2) CO2
Chlorine (Cl2) Turns damp blue litmus paper red, then bleaches it
Hydrogen (H2) “pops” with a lighted splint

Oxygen (O2) Relights a glowing splint

Sulfur dioxide Turns aqueous acidified potassium dichromate(VI) from
(SO2) orange to green
 QA
Redox reactions
Oxidising Acidified potassium manganate (VII) Purple solution turns colourless
agent (KMnO4)
Acidified potassium dichromate (VI) Orange solution turns green
(K2Cr2O7)
Halogen Solution turns darker Solution reduces to
Chlorine (Cl2) form Cl-
(due to formation Br2 but
or I2) when reacted oxidises Br- to Br2
with iodide/ bromide oxidises I- to I2
Reducing Potassium iodide (KI) Solution turns brown
agent Reactive metals (Al, Zn) Less reactive metal Displacement of
produced less reactive
metals
Hydrogen (H2) Reddish-brown solid (Cu) is Reduce CuO to Cu
formed
 QA y r u here
Colour characteristics of chemicals
Colour Inferences
Colourless Dilute acids, alkalis and solutions of salts of Group I, II and III metals
White Solid salts of Na+, K+, NH4+, Ca2+, Zn2+, Pb2+, Al3+
(amphoteric)
(Zn2+ salts colourless in water)
Black CuO, CuS, CoO, FeO, FeS, PbS, MnO2, I2 crystals
Grey Metals in powder form
Dark green Chromium salts
Light green Iron(II) and copper(II) salts
Blue/bluish green Hydrated copper(II) salts
Yellow/brown Solutions of iron(III) salts, PbI2, AgI
Pale pink Manganese(II) salts (Mn2+)
Purple KMnO4
QA
Precipitation reaction: formation of an insoluble solid
(precipitate) when two solutions are mixed together

Acid-base neutralisation reaction:

- acid: dissociates to produce hydrogen ions when
dissolved in water
- base: reacts with acids to produce a salt and
water; can be metal oxides or hydroxides
 Metals Reactivity: (Most to least)
K
QA Na Pot so can make a carbon zoo in tiny lead
Ca hydra copper silver gold platinum
Mg
Al Mine is backwards (most likely to be discharged on
top) (Because the less reactive, the more likely to
C discharge):
Z Stop chewing he lied the infant zebra ate
Fe my chocolate soup pie
Sn - Tin
(My friend’s mnemonic)
Pb please stop calling me a crazy zebra; i totally love his cheek; so good; pen
H
Cu
Ag
Au
Pt - Platinum
Thermal stability
- All metal oxides stable, except Ag, Au and
mercury
- All hydroxides except K and Na
- Gives water and metal oxide
- Ag, Au and mercury don’t exist in hydroxide form
- All carbonates, except K and Na
- Gives CO2 and metal oxide
- Ag, Au and mercury do not exist in carbonate form
Redox
Oxidation definitions (original/old definitions)
● Definition 1: oxygen
○ Oxidation → gain of oxygen
○ Reduction → loss of oxygen

● Definition 2: hydrogen
○ Oxidation → loss of hydrogen
○ Reduction → gain of hydrogen
Redox
Oxidation definition (modern definitions)
● Definition 3:electrons
○ Oxidation → loss of electrons
○ Reduction → gain of elections
○ LEO the lion says GER
○ Less Electrons Oxidise Gain Electrons Reduce

● Definition 4: oxidation state

○ Oxidation → increase in oxidation number
○ Reduction → decrease in oxidation number
 Redox
For the oxidation number thing
● Each atom in a pure element has an ox.no. of 0
● Monoatomic ion, ox.no. equal to charge on ion
(e.g. ox.no. of Mg2+ is +2)
● Oxygen: usually -2 except in peroxides that
contain O22- ion (e.g. H2O2) each oxygen ox.no.
is -1
Redox
For the oxidation number thing
● Ox.no. of Fluorine is always -1
● Sum of ox.no.s of all atoms in neutral
compound is 0
● Sum of ox.no.s in polyatomic ion equals
charge of ion
Redox - stuff idk where to put
● Disproportionation reaction
○ Same substance undergoes BOTH oxidation and
reduction
○ “When a substance is both oxidised and reduced
simultaneously to form two different products”
Also in some reactions when you add heat it’s
written like:
heat
___(_) + ___(_) ____(_) + ____(_)
Redox - metals reactivity thing (wait why is
this part of redox idk is it cos all the metals
are reduced in these reactions what is
chemistry
Highly reactive metals→ highly unreactive compounds
Cos its harder for less reactive metals to replace more reactive metals in reactions

● Metals always react by losing electrons (oxidising) to

form positive ions
● Greater tendency to lose electrons
○ → more reactive
○ → stronger reducing agents
Metal Observations on Reactions with...
Cold Water Steam Hcl

Potassium (K) Reacts very violently to produce KOH and (same as with cold water? Or should it be like Explosive reaction
H2. super super violent bomb explosive shit)
H2 is burned with a lilac flame
Sodium Reacts violently to produce NaOH and H2. (i mean if Mg is alr quite violent wouldn’t these Explosive reaction
H2 may catch fire and burn with a yellow be more intense)
flame.
Calcium Reacts readily to produce Ca(OH)2 and H2. (i tried googling but i only found like one source Reacts violently to produce H2.
that says it’s the same reaction as with water)
Magnesium Reacts very slowly to produce Mg(OH)2 and ~Hot~ Mg reacts violently to form MgO (white Reacts rapidly to produce H2.
H2. powder) and H2.
A bright white glow is produced during reaction.
Poor aluminium the notes Nah Sometimes Yesss it does react
didn’t mention you T_T :””( (explanation (not that important but still)) Reacts readily(??) to produce H2.
[Although Al reacts readily with steam to give
[Upon being exposed to air, aluminium Al2O3 and H2(g), the reaction does not always
instantly develops a layer of aluminium occur. This is due to a thin but strong layer of
oxide that is extremely unreactive protects aluminium oxide being coated onto the metal,
the metal from further oxidation.] thus preventing it from the reaction.]
Zinc Nah ~Hot~ Zn reacts readily to form ZnO and H2. ZnO Reacts moderately fast to produce
is yellow when hot and white when cold. H2.
Iron Nah ~Red-hot~ iron reacts slowly to form Reacts slowly to produce H2.
FeO/Fe2O3/Fe3O4 (idk it’s kinda dubious) and H2.
(basically everything else) Nah Nah Nah
Tin
Lead
Copper
Silver
Gold
Platinium
Metal oxides reactions
Reactions with
Metal Oxide
Carbon Hydrogen (metal oxides have to be heated)

K2O Nah Nah

Na2O
CaO
MgO (these metals are more reactive than carbon & hydrogen→ stronger reducing agent)
Al2O3
ZnO Oxide is reduced by carbon to form metal & CO. Heated metal oxides are reduced with hydrogen to form
FeO ZnO(s) + C(s) → Zn(s) + CO(g) metal and steam.
SnO2 FeO(s) + C(s) → Fe(s) + CO(g)
PbO SnO2(s) + 2C(s) → Sn(s) + 2CO(g) Dunno why even tho Zn, Fe & Pb are more reactive than
Pb(s) + C(s) → Pb(s) + CO(g) hydrogen they still react
CuO Oxide is reduced by carbon to form metal and CO2.
2CuO(s) + C(s) → 2Cu(s) + CO2(g)
Ag2O Oxide is reduced by heating without the need of a
reducing reagent.
2Ag2O(s) → 4 Ag(s) + O2(g)
Metal moosh metal oxide
More reactive metal higher tendency to form
positive ions compared to less reactive metal
→ more reactive metal reduces oxide of less
reactive metal
More reactive metal oxide + less reactive metal → no reaction
[____] does not react with [_____] as [_____] is less reactive than [____] and
hence cannot replace [_____] from it’s oxide.
Corrosion of metalsss
● Corrosion of Fe
○ In Iron the rust (corroded iron) is brittle and flaky and flakes away → new metal surface
exposed → corrosion → IT ALL RUSTS AWAY O_O
○ OXIDISES to form hydrated iron (III) oxide
● Ways to protect
○ Coat with paint/grease
○ Cover with plastic
○ Galvanising (zinc-plating) (usually iron/steel)
■ Electroplated by dipping into molten Zn/Sn
■ → thin film of Zn/Sn covers iron/steel
■ → prevents water & air from coming into contact with iron/steel surface
○ SACRIFICIAL PROTECTION (it exists trust me)
■ Put more reactive metal
■ → it sacrifices to protect iron/steel (it’s legit in the notes)
● [e.g. Mg rods attached to underground Fe pipe
● Mg more reactive than Fe → corrodes more easily
● Protects Fe from rusting as it oxidies more easily than Fe
● Mg donates its electrons to prevent it from rusting; as iron oxidises it will
immediately be reduced back to iron.]
Extraction of Metals
Metal Method of extraction

K Use of electricity to decompose

Na molten metal compounds
Ca
Mg
Increasing Al
reactivity
Zn Reducing metal oxides using carbon
Fe (refer to metal oxides reactions)
Pb
Cu
Ag

Au Naturally found uncombined as

metals
 Blast Furnace - Extraction of metal
● Carbon in coke is burned in a blast of hot air to produce CO2. This reaction
produces a lot of heat.
○ [C(s) + O2(g) ---heat---> CO2(g)]
● CO2 oxidises C in coke to form CO at high temperature of 1200ºC-1300ºC.
○ [CO2 (g)+ C(s) ---heat---> 2CO(g)]
● CO then reduces the [metal oxide] in the ore to [metal] and produces CO2.
○ [CO(g) + metal oxide ---heat---> metal + CO2(g)]
● The [metal] formed runs to bottom of furnace ← It’s extracted! :D
● Hot waste gases containing CO, CO2 and N2 escape through the top of furnace

● Limestone (Calcium carbonate) is decomposed by heat to produce CO2 and

calcium oxide (quicklime). [this reaction is used for see bottom]
○ [CaCO3(s) ---heat---> CaO(s) + CO2(g)]
● Iron ore contains other impurities (sand&clay), silicon oxides
● CaO (from 2nd reaction) reacts with acidic silicon (IV) oxide and other impurities to
form molten slag (Calcium silicate) [CaO(s) + SiO2(s) ---heat---> CaSiO3(l)]
Oxidising Agents
- Highly electronegative
- High oxidation number
- Are reduced
life
 sis
oly
ct r
Ele
Electrolysis
● Takes place in an electrolytic cell
● Conduction of electricity by an ionic compound (electrolyte)
when molten/dissolved in water, leading to the decomposition
of the electrolyte.
● Decomposition & Redox reaction
● Anode = oxidation
○ (Think A+ = gives a +ve charge)
○ Positively charged, attracts anions (hohoho all them a’s)
● Cathode = reduction
○ Negatively charged, attracts cations
Electrolysis
● Evidence that ions are held in lattice when
solid but free to move when (l) or (aq)
● Produces chemical change in non-
spontaneous reaction to occur
Drawing
- Must have:
- External power
source
- Anode and cathode
- Electrolyte
!! Electrolysis of WATER (memorise this!)
Yep and also take note that in
dilute aqueous solutions these are

2 H2O (l) → 2H2 (g) + O2 (g) usually the dominant reactions (see
selective discharge)
(cos hydrogen is relatively low
Oxidation half-reaction: reactivity → discharge more easily)

● 4OH- (aq) → O2 (g) + 2H2O (l) + 4 e-

(I just realised the letters make OHOHO e)

Reduction half-reaction:
● 2H+ (aq) + 2 e- → H2 (g)
Often sulfuric acid or nitric acid is added as pure water is a poor conductor
of electricity.
Electrolysis of molten ionic
compounds
● Liquid state
● Can conduct electricity due to mobile ions
● E.g. Pb(II)Br2
● At anode, oxidation occurs, e- removed
○ 2Br- (l) → Br2 (g) + 2e- *2Br and not Br 2 because the bromine
particles are ions when molten, not compounds)
○ Reddish brown and pungent gas evolved at anode
● At cathode, reduction occurs, e- added
○ Pb2+ (l) + 2e- → Pb (l)
○ Silvery liquid collected at the bottom of reaction vessel
Electrolysis of molten ionic
compounds
● Binary ionic compounds = contain ions from
only 2 different elements
Selective discharge
● When there are more than 1 type of cation
and anion present in the solution
● Depends on:
1. Position in the electrochemical series
2. Nature of electrode
3. Concentration of anion
Selective discharge
● When inert electrodes are used
○ Do not react. (Platinum, graphite)
1. Electrochemical series
a. Cations from elements higher in the electrochemical
series are LESS easily discharged. (lol remember the
mnemonic)
b. Anions
i. OH- is most ready to give up electrons
ii. (More) I > Br > Cl (less) likely. Grp 7, high En
iii. SULFATES and NITRATES do NOT discharge
Selective discharge (kinda a summary of the last few points in
the previous slide)

Cations Anions
K+
Na+
SO42-
NO3- } They don’t
discharge

Cl-
… Br-
... I-
Increasing
(refer to reactivity) ease of
discharge OH-
(mnemonic: So no clams bro? I OH)
Selective discharge
2. Nature of electrodes
a. Inert electrodes provide surface for electron transfer
b. Uh
c. If electrode is reactive electrode (low reactivity = discharge
easily) (uhhh the thing about Cu electrodes one side
will oxidise one side will reduce)
3. Anion concentration ONLY
a. Greater concentration will be preferentially
discharged, even if it’s higher in the EC series
Selective discharge
E.g. Electrolysis of NaOH (aq)
● Ions present:
○ Cathode: Na+, H+
○ Anode: OH-, Cl-
● Products formed: (see electrolysis of water)
○ H2 (g) and O2 (g) because
■ H is lower in the EC series
■ OH- is more easily discharged than Cl-
Real life examples
Extraction of Al
Purification of Copper
Need to elab
Electroplating
● Thin layer of metal is coated onto another
object
● Protect against corrosion
● And to improve appearance
● E.g. the gold orchid
○ The electrodes are plated (orchid)
■ Anode - plating metal
■ Cathode - orchid
○ Electrolytes need to contain ions of plating material
(gold)
Simple cells
● 2 diff metals in a single electrolyte
● Redox occurs tgt to cause a flow of e-
● From a more reactive metal to less
reactive metal → electrical energy
● Electrical energy is produced
● The further apart 2 metals are in the
reactivity series, the greater cell voltage is
produced (potential difference)
Simple cells
● More reactive metal becomes negative
electrode
● Less reactive metal becomes positive
electrode
● Electrons are transferred through the wire.
As a result, current is produced and reflected
on the
Simple Cells
E.g.
● Since Mg is higher in
the EC series, Mg is
the negative electrode.
In the voltaic cell, the -
ve electrode is the
anode. (reverse of electrolytic
cell)
Corrections
What is the function of the magnesium sulfate
solution?
- To act as an electrolyte (electricity passes
through) / To allow migration of ions /
Maintain electrical neutrality / Prevent build
up near electrodes.
NOT “to allow electrons to flow through”
Ene
rget
ics
Energetics
Change of enthalpy in a system a.k.a heat of
reaction is equal to amt of heat given off or
absorbed during a chemical or physical change
at constant pressure.

Measured in kJ/mol
Energetics
Breaking of bonds absorb energy
Creation of bonds release energy

It takes energy to break up with you.

Endothermic reaction
Heat is absorbed into the system, causing
surrounding temp to drop.
Products have higher energy content than
reactants.
Enthalpy is positive.
Endothermic reaction
Examples:
● Photosynthesis!
○ 6CO2(g) + 6H2O(l) → C6H12O6(s) + 6O2(g)
● Dissolution of ammonium salts in water
● Photodecomposition of silver halides
● Decomposition of CFCs in sunlight
● Nitrogen Oxide formation in car engine
Endothermic reaction ans
● Is the reaction endo/exothermic?
○ The reaction is endothermic as more energy was
required to break the bonds in the reactants
(name) than the energy released by making the
bonds in the product (name).
○ The energy content of the products is higher than
that of the reactants.
○ Energy was absorbed from the surroundings in the
form of heat, causing temp to drop.
Endothermic energy profile diagram
Exothermic reaction
Occurs when energy in the form of heat is
released, causing a temperature rise in the
surroundings.
Final energy content of products is lower that
reactant.
Enthalpy is negative.
Exothermic reaction
E.g.
● Combustion of fuels
● respiration
● neutralisation of acids and alkali
● displacement reactions
Exothermic reaction ans
● Is the reaction endo/exothermic?
○ The reaction was exothermic as more energy was
released to form new bonds in the product (name)
than the energy absorbed to break bonds in the
reactants (name).
○ Energy in the form of heat was given off to the
surroundings, causing a rise in surrounding temp.
○ The energy content of the products are lower than
that of the reactants.
Exothermic energy profile diagram
Bond energy/enthalpies
- Energy required to break the covalent bond
in one mol of gaseous molecules into its
atoms.
- kJ/mol
Hydrogen Fuel Cell
Reaction at cathode:
O2(g) + H2O (l) + e- → 4OH- (l)
Reaction at anode
2H2 (g) + 4OH- (l) → 4H2O(l) + 4e-
Overall reaction
2H2(g) + O2(g) → H2O(l)
Hydrogen Fuel Cell (pls know this)
Advantages
● Produce electricity indefinitely
● Converts chemical energy into electrical
electricity efficiently so that large amounts
of energy can be produced without loss of
energy
● Water is the only product. No pollutants (e.g.
soot, oxides of C or N, sulfur dioxides).