TO IMPROVE THE ACCURACY OF RESULTS IN AN EXPERIMENT:
1- Repeat each reading three times and obtain
the average (more accurate) result.
2 - Measure volumes of liquids using burette
instead of measuring cylinder
because it is more accurate.
3 - Measure volumes of gases given off in an experiment
using gas syringe because it is more accurate.
4 - When drying a solid:
- dry between filter papers
- or place it in a desiccator containing a drying agent such as anhydrous
calcium chloride.
Do not heat the crystals in an oven to dry it because
this may cause decomposition of the substance or loss of water of crystallization.
HAZARD LABELS
When drawing a graph:
not
In the tests on solids/solutions:
 If a solid is colored, it has a transition metal.
 If no transition metal, it is a white solid or colorless solution
 When solid is heated and condensation forms at the top of the
tube, it is a hydrated solid.
COLLECTING GASES
If gas is insoluble in water (ex. Oxygen, Hydrogen):
If gas is lighter than air (ex. Ammonia)
If gas is heavier than air (ex. Carbon dioxide)
To determine the volume of the gas
Delivery tube must be removed before heating is stopped
to prevent back suction of the water into the hot tube
which would break it.
Cotton wool may be used to:
- hold liquids
- prevent solids from passing through the
apparatus
- prevent loss of liquid by splashing
A suction pump is used in some experiments to
suck gases through the apparatus.
An airlock:
An airlock is used in experiments such as fermentation to allow gases
(ex. carbon dioxide) to escape but prevent air from entering.
In an experiment
The temperature may increase or decrease during the experiment
(exothermic/endothermic reactions).
However, after sometime, the temperature of the experiment will
return back to the original room temperature since reaction has
finished.
When doing an experiment involving measurement of
temperature:
To read the stopwatch:
Note the meaning of the following words:
reliable reading can be repeated to give the same results.
valid reading measures what is intended
accurate reading reading is near to the true reading. (Accurate
instrument was used.)
Precision same reading is obtained under the same conditions.
anomalous results is a measurement that does not fit the curve.
CHANGE OF STATE
Is a substance solid, liquid or gas at room temperature?
SEPARATING MIXTURES
1. FILTRATION: To separate an insoluble solid from a
liquid
Crystallization
To obtain a soluble salt from an aqueous solution:
Heat to point of crystallization.
Cool to form crystals. Filter
through filter paper and funnel.
Dry the crystals between
filter papers.
To know the point of crystallization:
Insert a glass rod into the hot solution and remove it. Crystals should form on
the tip.
A saturated solution is one in which no more solid can dissolve at a certain
temperature.
Investigation:
Beach sand is a mixture of sand and broken shells made of calcium carbonate.
Calcium carbonate reacts with dilute hydrochloric acid to form a solution of calcium
chloride. Plan an investigation to find out the percentage of shell material in a given
sample of beach sand.
Weigh 10 g of beach sand using a balance. Put into a beaker. Add
excess dil. HCl and stir with a glass rod. Calcium carbonate dissolves
and bubbles of carbon dioxide gas are given off. Filter through filter
paper and funnel. Sand is collected as residue. Dry the residue
between filter papers. Weigh. Subtract from original mass and
calculate percentage.
Investigation:
Ethanedioic acid dihydrate is a white crystalline solid. This acid is
water-soluble and is found in rhubarb leaves. Plan an investigation to
obtain crystals of ethanedioic acid from some rhubarb leaves. You
are provided with common laboratory apparatus, water and sand
Crush the leaves with some sand in a mortar and pestle. Put into a
beaker and add water. Stir with a glass rod. Filter through filter
paper and funnel. Heat the filtrate to point of crystallization. Cool.
Filter the crystals.
Investigation:
Copper (II) oxide and carbon are both black solids. Copper (II)
oxide reacts with dilute sulfuric acid to form aqueous copper (II)
sulfate. Carbon does not react with dilute sulfuric acid. You are
given a mixture of copper (II) oxide and carbon and access to dilute
sulfuric acid. Plan an experiment to investigate the percentage of
copper (II) oxide in the mixture.
Weigh 10 g of the mixture using a balance. Put the mixture into a
beaker. Add excess sulfuric acid and stir with a glass rod. Filter
through filter paper and funnel. Wash the residue with a few drops
of distilled water. Dry the residue between filter papers and weigh
using a balance. Subtract from the original mass and calculate the
percentage.
Simple Distillation
Fractional Distillation
Paper Chromatography
Rf value of a spot = distance traveled by the spot = X
Distance traveled by the solvent Y
Investigation:
An orange drink may contain artificial colours E110 (sunset yellow)
or E129 (Allura red). Plan an investigation to determine the
presence of these artificial colours in a sample of orange drink.
Put a spot of the orange drink next to a spot
of E110 and E129 on a baseline near the
bottom of a rectangular filter paper. Put the
paper into a beaker containing a very small
amount of water (below the baseline). Allow
the water to move up the paper. If the
orange drink has spots the same height as
the artificial colorings, they are present.
INDICATOR
It is a substance that has different colours in acids and
bases
Examples:
• Litmus paper
• Universal indicator paper
• Phenolphthalein
• Methyl orange
EFFECT ON LITMUS PAPER
• Acids
turn blue litmus to red
• Bases turn red litmus to blue.
UNIVERSAL INDICATOR PAPER
Phenolphthalein
Methyl Orange
Reactions of Acids and Bases
• Acid
+ Base
HCl + NaOH
pH2
pH12
• Acid
+ metal
Zn + 2 HCl
salt + water
NaCl + H2O
pH7
salt + hydrogen
ZnCl2 + H2
Acid + carbonate
Base + ammonium salt
salt + carbon dioxide + water
ammonia gas
PREPARATION OF SALTS
The method used depends on the type of salt and its
solubility in water
Soluble (aq)
 All nitrates
 Sodium, potassium and
ammonium salts
 All acids
 Barium chloride
 Copper sulphate
Insoluble (s)







Barium carbonate
Calcium carbonate
Silver chloride /bromide/iodide
Lead chloride /bromide/iodide
Barium sulphate
Copper oxide
All metals
To Prepare a salt, one of the following methods can be
used:
1- Titration:
To prepare a soluble salt from soluble reactants.
2- Neutralization:
To prepare a soluble salt from an insoluble reactant
3- Precipitation: To prepare an insoluble salt.
TITRATION:
soluble reactants
NaOH + HCl
soluble products
NaCl + H2O
To prepare Sodium chloride from sodium
hydroxide:
Place 25 cm3 NaOH solution in a flask using
a pipette. Add 3 drops phenolphthalein. The
solution turns pink. Add dil. HCl from a
burette until the pink color turns colourless.
Note the amount of acid used.
Repeat by adding the required amount of
acid to 25 cm3 of NaOH without using
indicator. Heat the solution to point of
crystallization. Cool. Filter the crystals.
Wash with a few drops of distilled water. Dry
between filter papers
Investigation:
Oven cleaners contain an aqueous solution of sodium hydroxide. Plan
an investigation to show which of two different oven cleaners contains
the more concentrated solution of sodium hydroxide. You are provided
with common laboratory apparatus and chemicals
Put 25 cm3 of the first NaOH solution into a flask using a pipette.
Add 3 drops phenolphthalein. The solution turns pink. Add dil.
HCl from a burette until the pink color disappears. Note the
volume of acid used. Repeat using 25 cm3 of the other alkali
solution. The one that uses more acid is more concentrated.
To determine which solution is more concentrated:
NEUTRALIZATION
Insoluble base + acid
CuO(s) + H2SO4
Add solid copper oxide (black
powder), using a spatula, to 25 cm3
sulphuric acid in a beaker until
excess solid remains in the beaker.
Filter through filter paper and
funnel. The excess copper oxide will
be collected as residue. Heat the
filtrate to point of crystallization.
Cool. Filter the crystals.
soluble salt + water
CuSO4 + H2O
PRECIPITATION:
AgNO3 + NaCl
soluble reactants
insoluble salt
AgCl (s) + NaNO3
Add silver nitrate solution to
sodium chloride solution in a
beaker. Silver chloride forms as
a white precipitate. Filter
through filter paper and funnel.
Wash the residue with a few
drops of distilled water. Dry the
residue between filter papers.
Anion
Carbonate
(CO3 2-)
Test
Add dil. HCl
Result/Observation
Effervescence, bubbles of gas
(carbon dioxide) that turn
limewater milky
Chloride
(Cl-)
Add dil. Nitric acid, then add White precipitate
aqueous silver nitrate
Bromide
(Br-)
Add dil. Nitric acid, then add Cream precipitate
aqueous silver nitrate
Iodide (I-)
Add dil. Nitric acid, then add Yellow precipitate
aqueous silver nitrate
Nitrate
(NO3-)
Add aqueous
Bubbles of gas (ammonia) are
sodium hydroxide produced that turn damp red
then aluminum foil: litmus paper to blue
warm gently
Sulphate
(SO4 2-)
Add dil. Nitric acid
then add aqueous
barium nitrate
White precipitate
Sulphite
(SO3 2-)
Add dilute
hydrochloric
acid, warm
gently
Bubbles of sulfur dioxide gas
produced which turn
acidified aqueous potassium
manganate(VII) from purple to
colourless
Add Aqueous
Sodium Hydroxide
Aluminium White precipitate,
soluble in excess
(Al3+)
giving a colorless
solution
Add Aqueous Ammonia
(Ammonium hydroxide
solution)
White precipitate, insoluble in
excess
Add Aqueous Sodium
Hydroxide
Ammonium
(NH4+)
(Aqueous
Ammonia
solution)
On warming, Bubbles
of ammonia gas are
given off that turn
damp red litmus
paper to blue.
Add Aqueous Sodium
Hydroxide
Calcium
(Ca2+)
White precipitate
insoluble in excess
Add Aqueous Ammonia
(Ammonium hydroxide
solution)
No precipitate or very slight
white precipitate
Add Aqueous Sodium
Hydroxide
Chromium
(Cr3+)
Green precipitate
soluble in excess
Add Aqueous Ammonia
(Ammonium hydroxide
solution)
Grey-green precipitate
insoluble in excess
Add Aqueous Sodium
Hydroxide
Copper
(Cu2+)
Add Aqueous Ammonia
(Ammonium hydroxide
solution)
Light blue
Light blue precipitate,
precipitate, insoluble soluble in excess giving a
dark blue solution
in excess
Add Aqueous Sodium
Hydroxide
Iron (II)
(Fe2+)
Add Aqueous Ammonia
(Ammonium hydroxide
solution)
Green precipitate, insoluble in excess
Add Aqueous Sodium
Hydroxide
Iron (III)
(Fe3+)
Add Aqueous Ammonia
(Ammonium hydroxide
solution)
Red-brown precipitate, insoluble in excess
Add Aqueous Sodium
Hydroxide
Zinc (Zn2+)
Add Aqueous Ammonia
(Ammonium hydroxide
solution)
White precipitate, soluble in excess giving a
colorless solution
TEST FOR GASES
Gas
Test
Ammonia Insert damp red litmus
(NH3)
paper
Result
Damp red litmus paper
turns blue
Gas
Test
Carbon dioxide Bubble the gas through
(CO2)
lime water
Result
Lime water turns
milky
Gas
Chlorine
(Cl2)
Test
Result
Insert damp blue litmus paper It bleaches
Gas
Hydrogen (H2)
Test
Insert a lighted splint
Result
It "pops"
Oxygen (O2)
Insert a glowing splint
It relights
Flame Tests
• Clean
a platinum wire by dipping it in conc. HCl.
• Dip the wire in the salt.
• Expose the wire to the non-luminous flame of the
bunsen burner.
• Observe the emitted light.
Flame Tests
Test for:
Water
(chemical
test)
Pure Water
(Physical
test)
Test
Result
Add anhydrous copper sulphate
It turns from white to blue
or
Add anhydrous cobalt chloride
paper
It turns from blue to pink
Heat the liquid to boiling
It should boil at 100 oC
Test for:
Alkenes
(C=C)
Test
Add Bromine water
Result
It turns from reddish
brown to colorless
Test for:
Test
Result
Ethanol Put a lighted splint near the The liquid
(organic liquid
catches
flammable
fire/burns
liquid)
Antiseptic smell
Ethanoic
acid
Appearance: colourless
liquid
Smell:
smell of vinegar
Flourine
Chlorine
Bromine
Iodine
Copper metal
Copper sulphate
Copper hydroxide
Copper carbonate
Copper oxide
Magnesium oxide
Nitrogen dioxide gas
Carbon
Manganese dioxide
yellow gas
green gas
Reddish brown liquid
Grey solid
reddish brown
blue
blue
green
black solid
white ash
reddish brown fumes
black solid
black solid (acts as catalyst )
STUDYING RATES OF REACTIONS
1- The volume of gas may be measured using a gas
syringe.
At the beginning: Reaction is fast because more particles are colliding.
At the end: Rate slows down and becomes zero because reactants are being
used up so less collisions.
Experiment (2) is faster than A; Experiment (1) is slower
Experiment (3) is faster using more reactants, so more gas is given off.
Increase in rate in experiment 2 may be due to:
- higher temperature
- higher concentration
- higher pressure
- use of better catalyst
- crushing the solid
2- Measuring change in mass
3- If a precipitate is formed in the Reaction
An aqueous solution of hydrogen peroxide decomposes very slowly to form
oxygen. The speed of decomposition can be increased by using a catalyst. Two
possible catalysts are the solids copper (II) oxide and chromium (III) oxide. Plan
an investigation to find out which of these two oxides is the better catalyst for
this decomposition.
Weigh 5 g of the first catalyst using a
balance. Put into a flask connected to
a gas syringe. Add 25 cm3 of
hydrogen peroxide solution to the
flask through a dropping funnel.
Determine the mass of gas collected
in two minutes. Repeat using 5 g of
the other catalyst and 25 cm3 of the
hydrogen peroxide solution. The one
that gives more gas in 2 minutes is the
better catalyst.
To prepare ethanol:
To obtain organic compounds from petroleum:
1-Fractional distillation Crude oil is heated and vaporized, then
cooled and condensed in a fractionating column to give different
fractions at different boiling points.
The one with less number of carbons (lower boiling point) is
collected at the top of the fractionating column.
2 - Cracking: Long chain hydrocarbons are broken down into
shorter chains using heat (400 – 700 oC) and catalyst (Aluminium
oxide Al2O3). This produces shorter chain hydrocarbons that are
more useful and have higher prices.
TO COMPARE ENERGY VALUE OF FUELS
Measure 30 cm3 of water into a test tube using
a measuring cylinder. Determine the initial
temperature of the water using a
thermometer.
Put 25 g of the first fuel into a spirit burner
using a balance. Place the burner below the
test tube and light it. Allow it to heat the
water for 5 minutes. Note the final
temperature of the water.
Repeat the experiment using 25 g of the other
fuel and the same volume of water for 5
minutes. Determine the final temperature of
the water. The fuel that causes a higher
increase in the temperature of the water is the
better fuel.
ELECTROLYSIS OF MOLTEN COMPOUNDS USING
INERT ELECTRODES:
ELECTROLYSIS OF AQUEOUS SOLUTIONS USING INERT
ELECTRODES
Purification of Copper:
Electroplating with silver:
To electroplate a spoon with silver:
Connect a silver rod to the positive end of the battery using copper
wires. Clean the surface of the spoon with sand paper.
Put the spoon in a beaker containing silver nitrate solution.
Connect the spoon to the negative end of the battery.
Ag
Ag+ + e-
Ag+ + e-
Ag
METALS:
REACTION OF SODIUM WITH COLD WATER:
Observations:
• Sodium floats, darts, melts.
• Vigorous fizzing due to evolution
of hydrogen gas.
• Bubbles
of hydrogen gas may catch fire.
• The solution formed turns litmus paper
to blue (alkaline).
REACTION OF CALCIUM WITH COLD WATER:
Observations:
• Calcium sinks
• Strong fizzing due to evolution
of hydrogen gas
• Solution becomes milky because
calcium hydroxide formed is not
soluble in water.
When Magnesium reacts with steam:
Observations:
Magnesium burns with a bright flame
White solid is formed.
Hydrogen gas is formed which burns.
When Zinc reacts with steam:
DISPLACEMENT REACTIONS:
Zn + CuSO4
ZnSO4 + Cu
Observations:
• Blue colour of Copper sulphate solution changes to colourless.
• Zinc dissolves
• Red
precipitate of copper is formed.
Weigh 5 g of the first metal using balance. Put into a conical flask
connected to a gas syringe. Add 25 cm3 of dil. HCl to the flask
using a dropping funnel. Determine the volume of gas collected in
5 minutes. Repeat using 5 g of each the other metals and same
volume of HCl. The one that gives more gas in 5 minutes is more
reactive.
Reaction of metals with oxygen to form metal
oxide.
Put 2 g of calcium into a crucible and lid and
weigh using balance. Cover the crucible and
heat it. Remove the cover several times to
allow more air in. Leave to cool. Weigh the
crucible and lid with its contents. Reheat and
reweigh several times till constant mass.
Subtract the initial mass of crucible + lid +
calcium from the final mass to obtain mass of
oxygen used.
REACTION OF METALS WITH DILUTE HCl:
Mg + 2 HCl
MgCl2 + H2
Zn + 2 HCl
ZnCl2 + H2
Cu + HCl
no reaction
CORROSION OF IRON
Iron reacts with both water and oxygen to form
hydrated Iron (III) oxide Fe2O3. xH2O reddish brown, flaky
Weigh iron nails separately using balance. Put into different test
tubes. Add 20 cm3 of one of the samples of water to each test tube.
Leave at room temperature for 5 days. Remove the nails. Put into
separate beakers. Dry in an oven. Weigh the nails. The one with
more increase in mass has rusted more.
To prevent rusting:
1- Painting
2- Coating with plastic
3- Coating with oil and grease
4- Galvanizing: covering the iron with a layer of
zinc by dipping or spraying.
5- Sacrificial protection: blocks of more reactive
metal are strapped to the steel.
6- Electroplating with silver or nickel or
chromium.
EXTRACTION OF ALUMINIUM
At the cathode (-)
Al 3+ + 3 eAl(l)
At the anode (+)
2 O2O2 + 4eC + O2
CO2
EXTRACTION OF IRON
Uses of Slag:
- In the Blast furnace: it floats on top of the Iron that is
formed to prevent its re-oxidation.
- When removed from the blast furnace: it is used in road
making.
Conversion of Iron into Steel:
1- Hot oxygen is bubbled through the
molten iron obtained from the blast
furnace to oxidize the impurities:
Carbon forms CO2 gas;
Sulphur forms SO2 gas;
and the gases escape
2- Calcium oxide (a base) is added to
react with the acidic impurities and
remove them in the form of slag.
3- To form different types of steel, the
required amount of carbon and other
metals (nickel, chromium) are then
added.
EXTRACTION OF COPPER FROM ITS CARBONATE
Copper carbonate (green solid) is placed in
an evaporating dish in a tube and heated to
give copper oxide (black solid). Carbon
dioxide gas evolves.
CuCO3
CuO + CO2
Hydrogen gas is then passed over the
heated copper (II) oxide in the tube and the
water vapour formed is condensed in a Utube placed in ice.
CuO + H2
Cu + H2O
The black copper (II) oxide changes to reddish
brown as copper is formed.
Extraction of copper from malachite (CuCO3 + insoluble
impurities):
Crush the malachite in a mortar and pestle. Add dil. HCl to the
solid malachite in a beaker. Copper carbonate solid dissolves to
form a blue solution of copper sulphate. Filter through filter paper
and funnel to remove insoluble impurities. Add magnesium
powder to the filtrate. A brown precipitate of copper is formed.
Filter through filter paper and funnel to obtain copper as residue.
Wash the residue with a few drops of distilled water and dry
between two filter papers.
Weigh 10 g of the first sample of coal using
balance. Put into an evaporating dish. Cover with
an inverted funnel connected to a delivery tube.
Put 20 cm3 of potassium dichromate into a test
tube using measuring cylinder. Put the delivery
tube into the test tube. Burn the coal and measure
the time taken for the orange colour to change to
green.
Repeat using 10 g of the other samples of coal
using same volume of fresh samples of potassium
dichromate. The one that causes faster change in
colour produces more sulphur dioxide.
Weigh 100 g of the fertilizer using balance. Put
into a beaker. Add 100 cm3 of water using
measuring cylinder. Heat and stir the solution
until no more solid dissolves. Leave to cool.
Filter through filter paper and funnel. Dry the
residue between filter papers. Weigh. Subtract
from original mass to determine the mass of
fertilizer that dissolved in the 100 cm3 of water.