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Chemistry Revision Notes
Qualitative Analysis
Cation tests
Fe2+ (e
...
FeSO4) – Pale green
NaOH: Dirty green ppt formed, insoluble in excess aq NaOH, ppt turned brown on standing
in air (Oxidizes to form Iron III Hydroxide)
NH3: Dirty green ppt formed, insoluble in excess aq NH3
Fe3+ (e
...
FeCl3) -‐ Yellow
NaOH: Reddish-‐brown ppt formed, insoluble in excess aq NaOH
NH3: Reddish-‐brown ppt formed, insoluble in excess aq NH3
Cu2+ (e
...
CuSO4) – Light blue
NaOH: Blue ppt formed, insoluble in excess aq NaOH
NH3: Blue ppt formed, soluble in excess aq NH3 to form a dark blue solution
Ca2+ (e
...
CaCl2) – Colourless
NaOH: White ppt formed, insoluble in excess aq NaOH
NH3: No visible reaction (or small quantity of CaCl2 formed is dissolved or no ppt
...
g
...
g
...
g
...
g
...
Therefore
the alkaline gas is ammonia
Na+ (e
...
NaCl) -‐ Colourless
NaOH: No ppt observed, no gas evolved
NH3: No ppt observed, no gas evolved
K+ (e
...
KNO3) -‐ Colourless
NaOH: No ppt observed, no gas evolved
NH3: No ppt observed, no gas evolved
Anion Tests
CO32-‐ (e
...
sodium carbonate)
• Add dilute HCl and warm gently – effervescence observed, gas is colourless and
odourless
o Bubble the gas produced into limewater – gas evolved forms a white
precipitate with limewater
o CO2 is evolved
• Add BaCl2 and dilute HCl OR Aq Ba(NO3)2 and dilute HNO3 – White ppt formed,
soluble in acids (reacts with acid), CO2 produced
• Add Aq AgNO3 and dilute HNO3 – White ppt formed, turned pale yellow, soluble in
dilute acid, CO2 produced
• Add aqueous Pb(NO3)2 – white ppt formed
SO32-‐ (sodium sulfite)
• Add dilute HCL and warm gently – A colourless and pungent gas evolved turns
o Test for gas using filter paper that has acidified potassium manganate (VII) -‐
Acidified potassium manganate (VII) paper turns colourless
o SO2 is evolved
• Add BaCl2 and dilute HCl OR Aq Ba(NO3)2 and dilute HNO3 – White ppt formed,
soluble in acids (reacts with acid), SO2 produced
• Add Aq AgNO3 and dilute HNO3 – White ppt formed, turned pale yellow, soluble in
dilute acid, CO2 produced
SO42-‐ (sodium sulfate)
• Add BaCl2 – white ppt formed
o Add an equal volume of dilute nitric acid – White ppt insoluble in dilute
nitric acid
• Add BaCl2 and dilute HCl OR Aq Ba(NO3)2 and dilute HNO3 – White ppt formed,
insoluble in acid
• Add aqueous Pb(NO3)2 – white ppt formed
Cl-‐ (sodium chloride)
• Add aqueous AgNO3 – white ppt formed
o Split mixture into two parts (A and B)
A
...
Add aq NH3 until a change is seen – white ppt soluble in aq NH3 to
form colourless solution
•
Add aqueous Pb(NO3)2 – White ppt formed, soluble in boiling water but reappeared
as white crystals on cooling
I-‐ (potassium iodide)
• Add aqueous Pb(NO3)2 – yellow ppt is formed
o Add an equal volume of dilute nitric acid – yellow ppt insoluble in dilute
HNO3
o Heat mixture till solid dissolves, cool under tap – Ppt is soluble on warming,
giving a colourless solution, golden yellow crystals form on cooling
• Add aqueous AgNO3 – Yellow ppt formed
o Split mixture into two parts (A and B)
A
...
Add aq NH3 until a change is seen – Yellow ppt insoluble in aq NH3
NO3-‐ (sodium nitrate)
• Add NaOH followed by Devarda’s alloy (or Al foil/powder)
o Heat gently and test gas evolved with red litmus paper – Vigorous
effervescence observed on warming
...
Add dilute nitric acid – Ppt insoluble in dilute HNO3
B
...
If
carbonate ions are present, carbon dioxide will be liberated
...
Acids and salts usually do not produce reactions except precipitation reaction or with
carbonates (liberating carbon dioxide) or sulfites (liberating sulfur dioxide)
Metal hydroxides like Al(OH)3, Zn(OH)2 and Pb(OH)2 are amphoteric in nature and can react
with both acids and bases
...
• Copper hydroxide (solid-‐blue ppt) + 4NH3 -‐> Tetra-‐amminecopper (II) ions,
Cu(NH3)42+ (deep blue solution) + 2OH-‐
• Zinc hydroxide (solid-‐white ppt) + 4NH3 -‐> Tetra-‐amminezinc (II) ions, Zn(NH3)42+
(colourless solution) + 2OH-‐
Flame tests
Different metals produce different coloured light as the quantum of energy absorbed for the
jump of electrons from one shell to the next differs and corresponds to a wavelength in the
visible light spectrum
• Barium – apple green
• Calcium – red
• Copper – green
• Lead – Blue/white
• Potassium – Lilac
• Sodium – Yellow/orange
Colour
Colour
Inferences
Colourless
Dilute acids, alkalis and solutions of Group I,
II and III
White
Solid salts of Na+, K+, NH4+, Ca2+, Zn2+,
Pb2+ and Al3+ (group I, II and III)
Black
Oxides (CuO, CuS, CoO, FeO, FeS, PbS,
MnO2, I2 crystals)
Dark green
Chromium salts
Light green
Iron (II) and copper (II) salts
Blue or bluish green
Hydrated copper (II) salts
Yellow or brown
Solutions of iron (III) salts (Fe3+), PbI2, AgI
Pale pink
Manganese (II) salts (Mn2+)
Purple
KMnO4
Tests for gases
Ammonia (only alkaline gas)
• Colourless and pungent
• Moist red litmus turns blue
• White fumes formed when gas rod dipped in concentrated HCl was brought near gas
Chlorine
• Pale yellowish-‐green and pungent
• Moist red litmus was bleached
• Moist blue litmus turned red and then bleached
Water vapour
• Colourless and odourless
• Blue cobalt chloride paper turned pink
Sulfur dioxide
• Colourless and pungent
• Orange acidified potassium dichromate turned green
• Purple acidified potassium manganate (VII) turned colourless
Carbon dioxide
• Colourless and odourless
• White ppt formed when gas is bubbled into limewater
• Soluble in excess gas forming a colourless solution
Oxygen
• Colourless and odourless gas evolved
• Glowing splint rekindled
Hydrogen
• Colourless and odourless gas evolved
• Lighted splint was extinguished with a pop sound
Thermal decomposition
Metal
Heat on Oxide
Heat on
Heat on
Hydroxide
carbonate
Potassium
Stable to heat
Stable to heat
Stable to heat
Sodium
Calcium
Magnesium
Aluminium
Zinc
Iron
Tin
Lead
Copper
Mercury
Silver
Gold
Decomposed to
the metal oxide
and water by
heat
Decomposed to
the metal
Decomposed to
the metal oxide
and carbon
dioxide by heat
Hydroxides do
not exist
Decomposed to
the metal,
carbon dioxide
and oxygen by
heat
Solubility Table
Soluble in water
Group 1 and ammonium salts
Most sulfates
Most chlorides (+ bromides and iodides,
Insoluble in water
-‐
PbSO4, CaSO4, BaSO4
PbCl2, AgCl
Heat on nitrate
Decomposed to
the metal nitrite
(NO2-‐) and
oxygen
Decomposed to
the metal oxide,
nitrogen dioxide
and oxygen by
heat
Decomposed to
the metal,
nitrogen dioxide
and oxygen by
heat
fluorides)
All nitrates
Group 1 and ammonium carbonates
Most carbonates
Group 1 and ammonium oxides,
Most oxides
CaO (sparingly soluble)
Group 1 and ammonium hydroxides
Most hydroxides
Ca(OH)2 (sparingly soluble)
Group 1 and ammonium phosphates
Most phosphates
Acids
Name
Chemical formula
Strong/weak
Organic/mineral
Hydrochloric acid
HCl
Strong
Mineral
Sulfuric acid
H2SO4
Strong
Mineral
Nitric acid
HNO3
Strong
Mineral
Ethanoic acid
CH3COOH
Weak
Organic
Carbonic acid
H2CO3
Weak
Mineral
Phosphoric acid
H3PO4
Weak
Mineral
Citric acid
C6H8O7
Weak
Organic
Acid + Base -‐> Salt + Water (neutralization)
• E
...
Sulphur acid + Sodium Hydroxide -‐> Sodium sulphate + Water
• Neutralization: H+ (aq) + OH-‐ (aq) -‐> H2O (aq)
Acid + Reactive Metal -‐> Salt + Hydrogen
• E
...
Hydrochloric Acid + Sodium -‐> Sodium Chloride + Hydrogen (Note: Violent
reaction)
• *The reactiveness of a metal decreases across the d-‐block and down a group
• Reactive metals: iron, zinc, magnesium, etc
...
g
...
g
...
g
...
g
...
g
...
g
...
g
...
g
...
g
...
g
...
g
...
g
...
g
...
g
...
g
...
After electrolysis, a metal product and a non-‐metal product are formed
...
)
• Half-‐reaction at cathode
o 2H+ (aq) + 2e-‐ -‐> H2 (g)
o 2H2O (l) + 2e-‐ -‐> H2 (g) + 2OH-‐ (aq)
o Hydrogen is produced (Effervescence observed, etc
...
The higher the ion in the series, the more difficult it is to discharge
2) The concentration of the anion in the electrolyte – the ion present in greater
concentration will be preferentially discharged
Electrochemical series
Cations (same as reactivity series)
Potassium (K) – The most difficult to discharge
Sodium (Na)
Calcium (Ca)
Magnesium (Mg)
Aluminium (Al)
Carbon (C)
Zinc (Zn)
Iron (Fe)
Tin (Sn)
Lead (Pb)
Hydrogen (H)
Copper (Cu)
Silver (Ag) – the easiest to discharge
Anions
SO42-‐ (will not discharge)
NO3-‐ (will not discharge)
Cl-‐
Br-‐
I-‐
OH-‐ -‐ the easiest to discharge
Note: Inert electrodes – electrodes that do not take part in any chemical reaction during
electrolysis
...
(E
...
Graphite or platinum electrodes)
Examples:
Electrolysis of aqueous sodium chloride
• At the cathode (Na+ and H+ are present)
o 2H+ (aq) + 2e-‐ -‐> H2(g)
o H+ is discharged instead of Na+ and hydrogen gas is produced
• At the anode (Cl-‐ and OH-‐ are present)
o 4OH-‐(aq) -‐> O2(g) + 2H2O(l) + 4e-‐
o OH-‐ is discharged instead of Cl-‐ and oxygen is produced
Electrolysis of concentrated aqueous sodium chloride
• At the cathode (Na+ and H+ are present)
o 2H+ (aq)+ 2e-‐ -‐> H2(g)
o H+ is discharged instead of Na+ and hydrogen gas is produced
o (Na+ is not discharged even though it’s present in a higher concentration
because it is way too stable than H+ to be discharged; -‐-‐-‐-‐-‐-‐-‐-‐-‐> cations are
unaffected by change in concentration as the difference in reactivity is too
great)
•
At the anode (Cl-‐ and OH-‐ are present)
• 2Cl-‐ (aq) -‐> Cl2 (g) + 2e-‐
• Cl-‐ is discharged instead of OH-‐ as it is concentrated and chlorine gas is
produced
Electrolysis of copper (II) sulfate solution using copper electrodes
• At the cathode (Cu+ and H+ are present)
o Cu2+ (aq) + 2e-‐ -‐> Cu(s)
o Copper metal is produced – reddish brown solid observed; at the same time,
copper deposits are forming on cathode so it is increasing in size
• At the anode (SO42-‐ and OH-‐ are present)
o Cu(s) -‐> Cu2+ (aq) + 2e-‐
o The anode is reacting to form copper ions, therefore the anode is decreasing
in size
The amount of copper (II) ions remains the same overall as for every one copper (II) ion
reduced in the cathode to make copper metal, a copper (II) ion is formed at the anode
...
) by electrolysis
• The ore is first converted into pure aluminium oxide (Al2O3), which is then dissolved
in molten cryolite (Na3AlF6) to lower the melting point of the aluminium oxide such
that electrolysis can be carried out
...
As they are more difficult to reduce than copper (II) ions, they will
remain in the electrolyte solution
The metal impurities less reactive than copper fall to the bottom of the reaction vessel as
anode sludge (they will not be oxidized)
The anode dissolves while purified copper (Cu2+ ions) deposits at the cathode
Electroplating
• A thin layer of a metal is used to ‘coat’ another object
• Protect objects again corrosion
• E
...
Plating an object with copper
Equation at anode (plating metal)-‐ Ag(s) -‐> Ag+ (aq) + e-‐
Equation at cathode (substance to be plated) -‐ Ag+ (aq) + e-‐ -‐> Ag(s)
Electrolyte solution – Contain ions of plating metal (e
...
AgNO3)
Electrodes – Cathode is the substance to be plated; anode is the plating metal
Simple cells
In a displacement reaction where the reactants are placed in contact with one another, the
electrons are transferred directly between each other the energy released is wasted as heat
to the surrounding
E
...
Copper (II) nitrate and zinc (redox/displacement reaction)
Zn(s) + Cu2+ (aq) -‐> Zn2+(aq) +Cu(s)
Oxidation half equation -‐> Zn(s) à Zn2+ (aq) + 2e-‐
Reduction half equation -‐> Cu2+(aq) + 2e-‐ à Cu(s)
Observations: Zinc ions are more reactive so they displace copper ions; copper metal is
formed as a reddish brown deposit; blue solution (copper II nitrate) fades as copper is
replaced by zinc and zinc nitrate is colourless
...
Electrochemical cell
• A set-‐up where a spontaneous redox reaction takes place and chemical energy is
converted to electrical energy
• Made up of two half-‐cells joined by a salt bridge
• Each half-‐cell consists of an electrode immersed in an electrolyte
• A salt bridge completes the circuit, while keeping the two electrolytes separate (it
can be either a glass tube filled with gel containing an inert electrolyte or a piece of
filter paper soaked in an inert electrolyte)
•
•
•
Oxidation takes place at the anode -‐> more reactive/less easily discharged metal
(negative terminal)
Reduction takes place at the cathode -‐> less reactive/more easily discharged metal
(positive terminal)
Electrons flow from the more reactive metal to the less reactive metal (anode to
cathode)
Simple cell
• Two different metals in a single electrolyte
• Electrons flow from the more reactive metal to the less reactive metal
• Anode (-‐) – more reactive (oxidation) – the electrons come from here
o More reactive -‐> more likely to oxidize -‐> it produces the electrons
• Cathode (+) – less reactive (reduction) – the electrons go here
The further apart the two metals are in the reactivity series, the greater the cell voltage
produced
...
When the reaction is
completed, the temperature decreases until it reaches room temperature
...
When the reaction is
completed, the temperature rises until it reaches room temperature
...
g
...
Bond energy or enthalpies
• Energy taken in to break a covalent bond
• In kJ/mol
Enthalpy change = Total energy absorbed for bond breaking (energy content of reactants)
– total energy released in bond-‐making (energy content of products)
Exothermic reaction – Final energy content of products is lower than energy content of
reactants
...
Endothermic reaction – Final energy content of products is greater than energy content of
reactants
...
How to calculate enthalpy change of a reaction (examples)
Q1) 2H2O(g) -‐> 2H2(g) + O2(g) – decomposition of water
O-‐H bond energy = 463
H-‐H bond energy = 436
O2 bond energy = 495
4(463) = 2(436)+146
Enthalpy change = 4(463) – 2(436) – 495 = +485kJ/mol
Therefore, this reaction is endothermic
...
GOOD LUCK FOR THE CHEMISTRY PAPER! :3