Chemistry practical ppt

KCSE CHEMISTRY PRACTICAL
REVISION GUIDE
FOR EXECELLENCE IN
CHEMISTRY PRACTICAL
2013 SERIES
COMPILED BY: ROBINSON CHEMISTRY TEACHER

TABLE OF CONTENT
• THE BASIC KNOWLEDGE FOR CHEMISTRY
PRACTICAL
• MOLE CONCEPT
• VOLUMETRIC ANALYSIS
• TITRATION
• PHYSICAL CHEMISTRY (ENERGY CHANGE,
REACTION RATES & SOLUBILITY)
• QUALITATIVE ANALYSIS (INORGANIC AND
ORGANIC)
• KCSE SAMPLE PAPERS WITH ANSWERS

BACKGROUND INFORMATION
Important skills required in chemistry practical
Practical is meant to assist the learners to learn
and develop the following skills;
• Manipulative skills
• Observation skills
• Reading skills
• Recording skills
• Computing skills
• Interpreting skills

• Manipulative skills- involve the correct and skillful handling of
the apparatus by the experimenter.
• Observation skills- involve the use of sense: Hearing, sight,
touch and smell to detect changes/reactions.
• Reading skills- refer to the ability to comprehend written or
printed information.
• Recording skills- entails writing down, for example,
measurements or statements of facts or other details for
reference.
• Computation skills- involve the use and application of
mathematical knowledge on the data collected.
• Interpretation skills- involve studying the data collected and
drawing conclusions based on the established chemistry
principles.

Areas to consider when preparing for
KCSE Chemistry
• Chemical symbols must be written correctly i.e.
• When the chemical symbol is just one letter, it
must be in capital.
• If two letters are used the first must be a capital
and the second one be a small letter.
• They must be separate (not joined).
• The small letter must not be bigger than the
capital ones.
• Must not be shaded.

•In the chemical formulae, any letter that represents one
atom of an element must be capital e.g. Na2CO3 and not
Na2Co3.
•A chemical equation MUST contain correct symbols and
formulae, balanced and state symbols to show the state of
mater of each of the substances involved.

Writing observations must be based on the
following:
• When there is colour change, both original
and final colours MUST be stated.
• Colours and smell of gases should be stated.
• Colours of precipitates should be stated.
• Colour changes for both blue and red litmus
papers MUST be stated when used.
• Colours of solutions must be stated.

A procedure must be correct to the letter e.g.
• Soluble salts must be prepared by adding EXCESS
metal to an acid, metal oxide to an acid or metal
carbonate to an acid.
• In separating a soluble solid from an insoluble
one, start by adding a convenient solvent to the
mixture. DO NOT say dissolve the mixture.
• Insoluble salts are prepared by precipitation
(mixing two solutions). Indicate what happens at
each step.

When naming or identifying substances, do it
as follows;
• When the questions leads identify, either the
name or formula is allowed.
• But when the question leads name, then just
give the exact name.
• Names of ions are not allowed only correct
formulae are accepted.
• If you give both the name and the formula,
both MUST be correct.

When plotting graphs, ensure they have the
following:
• Labeled axes.
• Appropriate scale.
• Plotted points must be visible.
• All points in the table must be plotted correctly.
• A convenient scale must make the graph occupy
at least ¾ page.
• Straight lines must be drawn with a ruler.
• Smooth curves must be drawn free hand.

Drawing diagrams
• Must be correctly drawn
• Apparatus should be proportional
• Labeling
• Must be workable i.e no lickage of the gas

Conditions under which the reactions were
carried out must be stated e.g.
• Catalyst
• Specific temperature or range of temperature
• Specific pressure
• Specific concentrations.
NB: Abbreviations are not accepted e.g. ppt
instead of precipitate

Filling the Burette:
• First rinse with about 5cm3 of the solution with
which it is to be filled.
• Discard the portion of solution used to rinse it.
• Fill the burette with the solution to the mark.
The level of the bottom of the meniscus is then
adjusted to be at the same level with the zero
mark. This is taken as the initial reading.
• Any convenient mark may be chosen as the initial
reading for any other subsequent filling

• Writing ions in words and wrong oxidation
states is not acceptable. Ions should be written
in their symbols and correct oxidation states e.g.
Zn2+ and not Zn+2.
• When mixing two solutions it is wrong to write:
• No reaction
• No change
• No chemical change/reaction
• Clear solution formed.

• Correct observation should be:
• A white precipitate formed.
• A colured (mention the colour) precipitate is
formed.
• No white (or no coloured) precipitate is
formed

• There is no white nor whitish solution. A
solution is colourless if no colour is seen in
the solution.
• Do not add-ish to colours(s) of precipitate of
solutions e.g. copper solution is blue not
bluish.

Using volumetric flask when:
a) Diluting solutions
• Measure the required volume of solution to be diluted and
transfer it to the flask.
• Add more distilled water until the bottom of the meniscus is
at the level of the mark on the neck.
b) Preparing solutions from solid substances.
• Dissolve the solid in a little water in a beaker.
• Transfer the solution into the volumetric flask.
• Add more distilled water until the bottom of the meniscus is
at the level of the mark on the neck of the volumetric flask.

VOLUMETRIC ANALYSIS
• This is a specialty in chemistry which deals
with measurements of volumes of solutions of
reactants, reacting them as you record the
reacted volumes and using the information to
carry out calculations based on the data
collected.

Terms associated with volumetric
analysis
• Standard solution:- is a solution of known
concentration which is used in titrimetry.
• Analyte:- is the test solution whose
concentration is to be determined.
• Equivalent volumes: are the volumes of both
the standard and analyte solutions which
react completely without any in excess. The
end of the reaction is known as the end point
or equivalent point.

Commonly used end point indicators
• For neutralization reactions, phenolphthalein
and methyl orange are used.
• Redox titrations do not require an indicator
since they change colour when reduced during
the reaction e.g.
• K2MnO4- (Potassium manganate (Vii) changes
colour from purple to colourless).
• K2Cr2O4- (Potassium dichromate (Vi) changes
colour from orange to green).

Indicator Colour in;
Acid Base Neutral solution
Methyl-orange
Screened methyl orange
Red/Pink
Purple
Yellow
Orange
Orange
Orange
Phenolphthalein Colourless Pink Colourless
Litmus solution Red Blue Purple
NB: Use 2 or 3 drops of indicator to get precise colour change/end point.

THE MOLE
• The mole is the chemist’s unit of
measurement of the amount of substances.
• It is a measuring unit just like a dozen, a pair,
or a gross e.g.
• A pair of shoes consists of 2 shoes.
• A dozen shoes consist of 12 shoes.
• A mole of shoes consist of 6.023 x 1023 shoes.

• In chemistry a mole of a substance is the
amount of that substance containing 6.023 x
1023 particles.
• In this case, particles can be atoms, ions,
molecules or electrons.
• The mole of a substance may be represented
by the chemical symbol or the chemical
formula of the substance referred to e.g.

• Cl- represents a mole of chlorine atoms.
• Cl2 – amole of chlorine molecules.
• O- a mole of oxygen atoms
• O2- a mole of oxygen molecules.
• Cu- amole of copper atoms.
• Cu+2- a mole of copper ions.
• e- a mole of electrons.

MOLAR MASS
• This refers to the mass of one mole of the
substances. The unit of measurement for the
mass is g/mol e.g

Amole of: Symbol or formula Molar mass (g/mol)
Oxygen atoms O 16
Oxygen molecules O2 32
Sodium chloride NaCl 58.5
Copper (ii) oxide CuO 79.5
Copper atoms Cu 63.5
Hydrogen Chloride
molecule
HCl 36.5

Concentration of a solution
• This is the amount of a solute dissolved to
make a given volume of solution.
• The concentration of a solution may be
expressed in terms of:
• mass per litre of solution (g/l) or (g/dm3)
• or
• moles per litre of solution (mol/l) or
(mol/dm3)

The number of moles in one litre solution is called
molarity (M)
A solution containing one mole of a substance per
litre of solution is referred to as a molar solution.
A molar solution, therefore, contains molar mass of
the solute per litre of the solution.

Relationship between Molarity,
Number of moles, Volume of solution,
RFM and Mass per litre of solution
• Molarity = Number of moles
Volume in litres
• Number of moles = Molarity x Volume inlitres
• Volume in litres = Number of moles
Molarity
Mass of one mole = RFM of the substance in grams

When diluting solutions use the
formula
• C1 V1 = C2 V2
Or
• M1V1=M2V2
Where;
• C1 is the initial concentration
• C2 is the final concentration
• V1 is the initial volume
• V2 is the final volume
• M1 is the initial molarity
• M2 is the final molarity

TITRATION
• Titration is a process in which a standard solution is carefully
reacted with an analyte with the sole aim of determining the
concentration. It involves measurement of volumes and
calculations.
• Tests on learner’s ability to manipulate the burette, pipette
and conical flask.
• The concentration determined can then be used to find out
either;
• Relative atomic mass
• Percentage composition in the compound used and
• Reacted mass
• Experiment is repeated 3 times until atleast two values with a
deviation of ±0.2 cm3 of each other are obtained (consistent
values).

Types of titration
• There are several types of titrations. But, only
four are commonly carried out at secondary
school level. These are:
• Direct titrations
• Back titrations
• Double indicator titrations
• Redox titrations.

Direct Titrations
These types involve the additions of a
standard solution to a fixed volume of the
analyte(solution whose concentration is not
known) with the aim of determining either;
• Its concentration;
• The RAM of one of the constituent elements
of the sample analyte; or
• The percentage composition of the
constituent elements.

Back Titration
• A back titration involves reacting an analyte
with excess of a reactant of known
concentration. The excess reactant is then
titrated against a standard solution.
• A back titration is preferred where a reaction
is too slow such that a sharp end point cannot
be attained easily or where a good end point
indicator cannot be identified.

Double indicator Titration
• Double indicator titration is carried out with the aim of
determining the composition of a mixture of different
compounds.
Example
• You are provided with a solution X that contains 12.72
g dm-3 of Na2CO3 mixed with an unknown mass of
NaHCO3 . 25cm3 of the solution mixture is first titrated
with 0.1 M HCl using phenolphthalein indicator. The
analyte is further titrated with the acid using methyl
orange indicator.

FOR A SUCCESSFUL VOLUMETRIC ANALYSIS
CONSIDER DOING THE FOLLOWING:
• Complete the table and ensure that your values are
consistent i.e. within ±0.2 of each other.
• A candidate may be penalized the mark(s) for complete
table if:
• Arithmetic’s are wrong
• Table is inverted i.e. values being misplaced e.g. initial
burette reading in the column of final burette reading and
vice versa.
• Values are above 50cm3 unless explained.
• Values are unrealistic e.g. above 100cm3 or below 0.5 cm3

• The candidate should use 1 0r 2 decimal
places when filling the table. Uniformity of
the decimal is a MUST.
• If you use 2 decimal places, the 2nd one must
be a zero or a five e.g. 20.10 or 20.15 and can
NEVER BE 20.13 or 20.16 etc.
• Values to be averages must be within ±0.2 of
each other.

• All working must be shown starting with the
calculations of the average volume (the principle
of averaging) all the way to moles and molarity.
• Transfer the values infact i.e. values as they are
from the preceding questions.
• Use correct units. Wrong units call for penalty of
½ mk. Better leave it if not sure
• It is important to know the colour of indicators in
acids and basic solutions.

PHYSICAL CHEMISTRY
This area deals with:
• Thermo-chemistry/energy changes
• Rates of reactions
• Solubility
• Electro-chemistry

Thermo-chemistry/energy changes
• Deals with energy in a system. This is because
as chemicals or physical changes take place, a
system may give out heat energy to the
surrounding or absorb from it.
• Any experimenter has to be very keen in
collecting accurate data and making correct
observations.
• Any one given chemical has a certain amount
of energy associated with it.

The amount of energy associated with a system is
referred to as the enthalpy of the system (symbol H)
The amount of energy gained or lost in known as
enthalpy change (symbol ΔH)
A reaction in which energy is gained is referred to as
an endothermic reaction.
A reaction in which energy is evolved /gained is said
to be an exothermic reaction.
NB: THESE TYPES OF CHANGES ARE NOTED SINCE
THEY INVOLVE TEMPERATURE RISE OR FALL ON THE
THERMOMETOR.

Energy level diagrams
An energy LEVEL diagram illustrates the; change in
energy, type of reaction, reactants and products in a
tabular form.
LEVEL A LEVEL B
H1
PRODUCTS
REACTANT
PRODUCTS
REACTANTS
H2
H2
H1

Where;
• H1 is the initial enthalpy
• H2 is the final enthalpy
• ΔH is the enthalpy change.
• In calculations, use the formula;
ΔH= H2 – H1

Which energy level diagram represents;
•An endothermic reaction?
•An exothermic reactions?
Answer:
•Level A
•Level B

Terms associated with enthalpy
changes
• Heat of Reaction (ΔH reaction): this refers to the
amount of heat energy absorbed or evolved
when a reactants change to products.
• Standard heat of Formation (ΔHf): is the amount
of heat energy absorbed or liberated when one
mole (1M) of a substance is formed from its
constituent elements under standard conditions
of temperature and pressure.
• Heat of Combustion (ΔHc): is the quantity of heat
evolved when a given mass of a chemical
substance is burnt in excess oxygen.

•Molar Heat of combustion (ΔHc): is the amount of
heat change when one mole of a pure substance is
burnt in excess oxygen under standard conditions of
temp and pressure.
•Latent Heat: this is the amount of heat energy
required to cause the physical change of a given mass
of a liquid to vapour at constant temperature i.e. heat
needed for; melting, freezing, evaporation, or
condensation.
•Heat of neutralization (Hneut): is the amount of heat
energy evolved when one mole of a solution of an
acid neutralizes one mole of alkali solution to form an
infinitely dilute solution.

•An infinite dilute solution is one which is so much
dilute such that there is no further enthalpy change
that occurs when more water is added.
•Lattice Energy (ΔHlatt): is the amount of energy
liberated when isolated gaseous ions combine to form
an ionic solid.
•Hydration Energy (ΔHhyd): is the amount of heat
energy liberated when isolated gaseous ions interact
with and are surrounded by water molecules.

Determination of Enthalpy changes.
Most reactions in thermo-chemistry occur in aqueous
solutions.
The following assumptions are always made when
working out enthalpy changes.
•The density of the solution is always equal to that of
water. Which is 1 x 103 kgcm-3
•The specific heat capacity is always the same as that
of water. Which is 4.2 KJKg-1K-1
•The solution formed is infinite dilute
•No heat is lost to the surrounding or gained from the
surrounding during the experiment.
•Dissolving a small amount of solid does not
significantly alter the final volume of the solution.

the sources of error and account for
the differences between
experimental and theoretical values.
• EXAMPLES
• Suppose excess iron is added to 25cm3 of a
0.2M CuSO4 solution at 25⁰c. if the highest
temperature reached is 35. Determine the
molar enthalpy of the reaction.
• It should be assumed that all the copper ions
in the 25cm3 of 0.2M CuSO4 are displaced by
iron.
• Equation for the reaction;

Equation for the reaction;
Cu2+(aq) + Fe(s) Cu (s) + Fe2+(aq)
Mole of copper displaced = Molarity x volume in litres
= 0.2 mol x 25cm3
1000cm3
= 0.005 moles
Mole ratio of Fe: Cu = 1:1
Therefore no. of moles of iron oxidized to Iron (II) ions = no.
of moles of copper (II) ions reduced to copper
= 0.005
moles
Amount of heat (Q) = Mass in Kg (M) x specific heat capacity
(c) x change in temp (ΔT)
Q = MC ΔT

= 25g x 4.2KJ x (35.2 -25) K
1000g x kgk
= 0.025 kg x 4.2kj x 10.2 K
Kg K
= 107.1 KJ
This means that 0.005 moles evolved 107.1 Kj.
Molar heat of solution is the amount of heat
evolved if one was used
Which is 107.1 KJ x 1 mole
0.005 mole
= 212 Kj/mol

Molar heat of solution is the amount of heat evolved
if one mole was used
Which is 107.1 KJ x 1 mole
0.005 mole
= 212 Kj/mol

Molar Heat of Neutralization of sodium Hydroxide.
This can be investigated by determining the change in
temperature that occur when different volumes of 2M
solution of unknown acid. Record the result in the table.
Sample results
Experiment I II III IV V VI VII
Volume of 2m NaOH (aq) 10 15 20 25 30 35
Volume of 2M Acid x Initial
temperature(⁰C)
35 30 25 20 15 10
Final temperature (⁰C) 24.0 24.0 24
.0
24.0 24.0 24.0
Temperature change (⁰C) 29.0 33.0 34
.5
37.5 36.5 33.0
5.0 9.0 10
.5
13.5 12.0 9.0

Questions
•Plot a graph of the rise in temperature against the volume of
sodium hydroxide solution added.
•From the graph determine the greatest temperature change
•What is the significance of the greatest temperature
change?
•Assuming that the density of the resulting solution in
experiments is 4.2jkg-1/k-1 calculate the molar heat of
neutralization of 2m NaoH (aq).

Experiment to investigate molar Enthalpy of Displacement.
•You are provided with:
•Solution Y containing 0.2 moles of Copper (II) Sulphate per
litre of solution
•Solid Z
•You are required to determine the heat evolved when one
mole of solid Y reacts with solid Z.
Procedure:
•Measure 40cm3 of solution Y and place it into an insolated
50cm3 plastic beaker.
•Stir the solution with the help of the thermometer and
record its temperature after every ½ minute for 1 ½ minutes.
•After exactly 2 minutes, add all the solid Z provided and
continue stirring the mixture while recording the temperature
of solution and complete the table 1 below.

Time
mins
½ 1 1 ½ 2 2 ½ 3 3 ½ 4 4 ½ 5 5 ½ 6 6 ½
Temp
•On the graph paper provided plot a graph of temperature
against time.
•From your graph, determine the maximum temperature
change.
•Given that the density of the solution is 1g/cm3, determine
the quantity of heat evolved when the 40cm3 of solution Y is
reacted completely with solid Z.
•Given that solid Z is Zinc powder, write an ionic equation of
the reaction which occurs.
•Determine the moles of copper ions used up in the reaction.

•Determine the amount of heat that would be evolved if one
mole of Copper (II) ions were used up.
•State with reasons whether the value obtained in this
reaction is higher or lower than the one obtained from data
books.

SOLUBILITY AND SOLUBILITY CURVES
• Solubility is the maximum amount of a solute
which can dissolve in a given mass of a solvent at
a given temperature.
• A water soluble substance dissolve in water until
no more can dissolve at a given temperature.
When this happens, the solution is said to be
saturated,
• The temperature at which a given mass of solvent
cannot dissolve any more solute is called
saturation temperature.

When such a solution is cooled below its saturation
temperature, the solution forms crystals of the solute.
Solubility of most solids increases with increase in
temperature; however, solubility of a few solids
decreases with temperature rise.
Solubility is expressed in or mole (mol.) per 100g of
the solvent.
Solubility of a substance can be determined
accurately by experiments in different ways as
illustrated below:

Rates of reactions
This is also referred to as chemical kinetics.
• It involves the investigation on the effects of changing one
of the conditions/parameters on the time taken for
reactants to be consumed or the products to be formed.
Factors affecting the rates of reactions include:
• Concentration of reactants
• Surface area (partial size)
• Pressure
• Temperature
• Presence or absence of a catalyst

The dependency of a given factor is investigated if the rest are kept
constant.
Determination of rates of reactions involves measurements of; time,
temperature, volume and the mass of reactants or products.
For good results the learner must be able to use calibrated instruments
accurately.
Plotting of graphs using data collected is a common feature in the study
of rates of reaction.
The following information and skills on graph work are a requirement
for the learner;
•Giving a graph a concise, accurate and specific title
•Choosing and using a suitable scale
•Labeling the axes accurately
•Plotting the collected data correctly
•Joining all plotted points for a curve or the best line of fit for a linear
graph

Qualitative Analysis
• Qualitative analysis refers to the process of carrying
out chemical tests on substances with the sole aim of
identifying them.
• Involves identification of ions without taking any
precise measurements.
• It involves close and careful observations that are
made on the chemical changes that occur during
chemical reactions.
• Reasonable conclusions/inferences are made based on
the observations e.g. when a substance is dissolved in
water, heated, acid is added, NaOH (aq) or NH3 (aq).

Qualitative analysis is divided into two parts:
• Part 1: Qualitative inorganic analysis
• Part 2: Qualitative organic analysis
• In both areas, to obtain good results, a
student must be able to:
• Accurately identify the various test reagents.
• Identify what these reagents tests.
• Predict the expected results.

The following are important points to
note
• Formation of precipitate generally leads to the
identification of cations.
• Evolution of a gas leads to the identification of anions.
• Testing for the presence of cations involves the use of dilute
NaOH or ammonia solutions as the preliminary test
reagents.
• Identification of anions involves the use of dilute HCl, (vi)
acid or nitric (v) acid as the preliminary test reagents.
• Flame tests involves the use of Nichrome wire or clean
metallic spatula or glass rods dipped in a solution
containing the metal ions and hating it in a non-luminous
flame.

• Colour of residues or oxides of certain metals is
also necessary in identifying ions e.g.
• Red resides when hot, yellow on cooling indicates
presence of Pb+2 as the residue is PbO.
• Black reside when hot or cold likely suggests FeO,
CuO hence Cu2+, Fe2+ ions might be present but
further test needed to eliminate one.
• Yellow reside when hot and white on cooling
indicates that the residue is ZnO, therefore Zn2+
ions present.

Inorganic Analysis
• The basic test for gaseous products include:
• Colour
• Smell
• Effect on limewater
• Effect on a glowing or burning splint.
• NB: Students often ignore basic tests for gases
even when one has reasons to believe a gas is
being evolved.

• Identification of a gas being evolved may be
one of the clues required to identify the
substance being tested.
Background information on salts.
• Solubility of salts
• Effects of heat on salts

Solubility of salts
• All nitrates are soluble in water
• All chlorides are soluble in water except AgCl and
PbCl2.Lead chloride dissolves on warming and
crystallizes on cooling.
• All carbonates are insoluble in water except
Na2CO3, K2CO3 and (NH4)2CO3.
• All sulphates are soluble in water except PbSO4,
BaSO4 and CaSO4. Calcium sulphate is slightly
soluble.
• All sulphites are insoluble in water except
ammonium and alkali metals sulphites.

• : Barium sulphite dissolves in dilute nitric (V)
acid and HCl while Barium sulphate does not.
This acts as a test for distinguishing the two
salts.
THE TABLES BELOW SHOWS THE PHYSICAL
NATURE AND EFFECT OF HEAT ON SALTS

Colour of salt Cation present
White solids & forms
colourless solutions
K+, Ca2+,Na+, Mg2+, Al3+,
Zn2+, Ba2+,Ag+, Pb2+,NH+
4
Pale green Fe2+
Green Ni2+ and Cr3+
Yellow/brown Fe3+
Blue/ pale green
depending on the
compound
Cu2+

Nitrate of: Effect of gentle
heating
Effect of strong
heating
K and Na Melts to a
colourless liquid
Decomposes to
yield oxygen and
nitrites
Ca, Mg, Al, Zn, Fe,
Pb and Cu
Melt to a liquid Decomposes to
produce oxygen,
nitrogen (Iv) oxide
and metal oxides.
NH4
+ Melts to a
colourless liquid
Decomposes to
nitrogen (i) oxide
and water

Carbonate of: Effect of gentle
heating:
Effect of strong
heating:
K and Na No effect Do not decompose
Ca,Mg,Zn, Pb, Fe and
Cu
No effect Decomposes to
produce carbon (iv)
oxide gas and metal
oxide.
NaHCO3 No effect Decomposes to form
carbon (iv) oxide,
water and metal
carbonate
(NH4)2CO3 Slow decomposition Decomposes to form
ammonia, carbon (iv)
oxide and water

Effect of heat on chlorides
Heat has no effect on chlorides.
Effect of heat on Sulphates :-
Iron (II) Sulphate decomposes on heating
liberating mixture of sulphur (IV) and sulphur
(VI) oxides and the oxide of the metals.
Ammonium sulphate decomposes to
ammonia gas and ammonium hydrogen
sulphate.

Gas Colour and
smell
Test Test result
Ammonia (NH3) - Colourless
- Has a pungent
smell
Hold moist litmus
paper ( or universal
indicator paper) in
the gas.
Indicator turns blue
Carbon (IV)
Oxide
- Colourless
- Odourless
Bubble the gas
through lime water
(calcium hydroxide
solution)
White precipitate of
calcium carbonate is
formed.
Chlorine (Cl2) - Pale green
- Has a choking
smell
Hold moist litmus
paper (or universal
indicator paper) in
the gas.
Indicator paper is
bleached. (Wet blue
litmus will turn red
first before getting
bleached.)
Properties and Tests for Common Gases

Hydrogen (H2) - Colourless
- Odourless
Hold a burning
wooden splint in
the gas.
Burns with a pop
sound.
Oxygen (O2) - Colourless
- Has a choking
smell
Hold a glowing
wooden splint in
gas
Relights the splint

• NB:
• Do not smell gases directly instead lift it
towards the nostrils
• Litmus paper used must be moist and should
not touch the walls or the test tube.
• Both blue and red litmus paper must be used.

Identification of cations
• Cations are positively charged ions. They are
identified through tests in which precipitates and
complex ions are formed when sodium hydroxide
or aqueous ammonia is added to their salt
solutions.
• The method is based on the principle that,
insoluble metallic hydroxide are formed when
two to three drops of aqueous ammonia or
sodium hydroxide are added to the test solution.

• NB: The colour of some of the precipitate and
solubility of some of the hydroxides in excess
alkali, help to identify the ions present in the
substance.

Cati
on
Reaction with NaOH
(aq)
Reaction with aqueous ammonia
(NH3)
Ba2+ - White precipitate
formed
- Precipitate insoluble
in excess alkali
- White precipitate formed
- Precipitate insoluble in excess
aqueous ammonia
Ca2+ - White precipitate
formed
in excess alkali
aqueous ammonia
Mg2+ - White precipitate
formed
in excess alkali
aqueous ammonia

Al3+ - White precipitate
formed
- Precipitate soluble in
excess alkali
- Precipitate insoluble in
excess aqueous ammonia
NH+
4 - No white precipitate
- When heated a gas
which turns moist red
litmus paper blue is
evolved
- No white precipitate
- When heated a gas which
turns moist red litmus
paper blue is evolved.
Zn2+ - White precipitate
formed
excess alkali

Fe2+ - Green precipitate formed
excess alkali
- Green precipitate formed
Fe3+ - Brown precipitate formed
excess alkali
- Brown precipitate formed
Pb2+ - White precipitate formed
excess alkali

Cu2+ - Blue precipitate
formed
- Precipitate
insoluble in excess
alkali
- Blue precipitate formed
excess aqueous
ammonia forming a
deep blue solution.

• NB: the precipitate dissolves due to the formation of complex ions
e.g.
• Amphoteric hydroxides of Aluminiu. Zinc and lead dissolves in
excess NaOH.
• Pb(OH)2 (s) + 2OH- (aq) Pb(OH)4
2- (aq) + 2Na+
• Zn(OH)2 (s) + 4NH3(aq) {Zn (NH3)4}2+(aq) + 2OH-(aq)
• Al(OH)3 (s) + OH- (aq) Al(OH)4
-(aq) + Na+
(aq)
• Zinc and copper hydroxide dissolve in excess ammonia solution to
form colourless tetra amine zincate and deep blue tetra amine
copper respectively.
• Zn(OH)2 (s) + 4NH3(aq) { Zn (NH3)4}2+(aq) + 2OH-(aq)
• Tetramine zinc (II) ion.
• Cu(OH)2 (s) + 4NH3(aq) { Cu(NH3) }2+(aq) + 2OH-(aq)
Tetramine Copper (II) ion

How to make conclusions/inferences
based on the observations
• If on adding few drops of NaOH or ammonia
solutions, white precipitate is formed, conclude
by writing;
• Either; Ca2+, Mg2+,Al3+, Zn2+,Pb2+ or Ba2+ ions
present.
• If green precipitate is formed, Fe2+ ions present.
• If brown precipitate is formed, Fe3+ ions present.
• If blue precipitate is formed, Cu2+ ions present.

•If on adding excess NaOH solution white
precipitate persisted (insoluble in excess NaOH),
conclude;
Ca2+, Mg2+ or Ba2+ ions present.
•If on adding excess NaOH solution white
precipitate dissolves (soluble in excess NaOH)
conclude.
Al3+, Zn2+ or Pb2+ ions present.
•If on adding excess ammonia solution white
precipitate persisted (insoluble in excess), conclude;
Either Ba2+, Ca2+, Mg2+, Al3+, or Pb2+ ions present.
•If on adding excess ammonia solution white
precipitate dissolves (soluble in excess) conclude;
Zn2+ ions confirmed present.

For cations which exhibit similar properties in both
dilute NaOH and aqueous ammonia. A confirmatory
test has to be performed to distinguish them e.g. Al3+
ions and Pb2+ ions which form white precipitates
which are soluble in excess NaOH but insoluble in
excess aqueous ammonia.
To distinguish them, two to three drops of either
sodium chloride (NaCl), sodium sulphate (NaSO4) or
potassium iodine KI are added. Lead ions form
insoluble chloride, sulphate or iodide respectively.

Apart from using NaOH and ammonia
solutions to identify cations flame test can
also be used.

Addition of dilute acids
Salt solutions reacts with dilute acids
especially HCl & H2SO4, hence are used to
confirm the preference of cations.

Test Observations Inferences
1.Add dilute HCl
to solution
- White ppt
formed
Pb2+,Hg2+ or Ag+
ions to be
present
- No white ppt is
formed
Pb2+, Mg2+ and
Ag+ absent
1.Add dilute
H2SO4
- White ppt
formed
Pb2+, Ca2+ or Ba2+
present
- No white ppt
formed
Pb2+, Ca2+ and
Ba2+ ions absent

Flame test for some cations
This method is based on the colours the cations
render a flame when heated.
Procedure.
•Dip a nichrome wire into concentrated HCl acid.
•Heat the end that has been dipped in the acid using
a flame until it cease to impart colour on the flame.
This is done to clean the wire.
•Dip the cleaned wire into the solution containing the
test ion and then heat it in the hottest part of the
flame. The test ion imparts its characteristics colour.
NB: platinum wire, metallic spatula or glass rod can
be used instead of the nichrome wire.

Cation Colour of the flame produced
Lithium (Li+) Crimson (deep red)
Sodium (Na+) Yellow
Potassium (K+) Lilac (Purple)
Calcium (Ca2+) Red
Copper (Cu2+) Blue-green
Barium (Ba2+) Apple green

CONFIRMATION TESTS FOR ANIONS
• Anions are negatively charged ions
• Most of the anions are radicals
• They include:- CO3
2-, HCO3
-, SO4
2- & NO3
-
• Examples of non-radical anions are S2- & Cl-,
Br-,I-

Anions Test Observations
SO4
2- Add barium
chloride and mix
followed by dilute
nitric or HCl acid
- A white
precipitate
insoluble in
dilute acid
SO3
2- Add Barium nitrate
(Ba(NO2)2 or
Barium Chloride
(BaCl2) and mix
followed by dilute
HNO3 or HCl
- A white
precipitate
soluble in dilute
acid

CO3
2-,
orHCO3
-
,
Heat the solid sample - K and Na carbonates
do not decompose all
other carbonates
decomposes on
heating
Add dilute acid - Effervescence of CO2
which from white ppt
in lime water is
evolved.
S2- Add dilute acid to solid
Test with lead acetate
- A gas is produced
- Gas blackens Lead (II)
acetate paper

Cl- Add AgNO3 or Pb (NO3)2
followed by dilute HNO3
- White precipitate
insoluble in HNO3
(aq)
- Ppt soluble in
excess aqueous
NH3 solution
NO3
- Add a small amount of freshly
prepared ferrous sulphate
solution.
Gently pour conc. H2SO4 down
the side of the ppt tube such
that the solutions do not mix
- A brown ring
may be formed
at the
interface/junctio
n of the two
solutions.

NB: preliminary and confirmatory Tests for
Chlorides, Nitrates, Sulphates and Sulphites.
• Confirmatory Tests for Nitrate Ions (NO3
-)
• When the nitrates of sodium potassium are
heated, the nitrates of the metals and oxygen
gas are produced.
• 2NaNO3 (s) 2NaNO2 (s) + O2 (g)
• If the test solid is heated in a closed boiling
tube, and a glowing splint is inserted at the
mouth of the tube, the splint relights.

• Nitrates of heavy metals decompose on heating to produce
brown nitrogen (IV) oxide gas, the respective metal oxide
and oxygen.
• 2Pb(NO3)2 2PbO (s) + 4NO2(g) + O2 (g)
The Brown Ring Test
• Freshly prepared solution of iron (II) sulphate to the test
solution, followed by concentrated sulphuric (VI) acid. If a
brown ring is formed at the interface of the two solutions,
this indicates the presence of nitrate ions in the test
solutions.

PRELIMINARY TESTS FOR ORGANIC
COMPOUND
• Organic compounds includes:
• Hydrocarbons
• Alkanols
• Alkanals etc
Flame test is used to test for hydrocarbons and
alkanals (alcohols).

Blue non- sooty flame shows presence of c c
Yellow smoky flame shows presence of c c
or
c c

TEST OBSERVATION INFERENCE
PH of the aqueous
solution
pH lesser than 7 Weak carboxylic acid pH
4 to 6.
O
(R C O H)
pH greater than 7 Compound is either an
alkanol R-OH, ester
O
(R C O R) or an
alkali metal salt of
carboxylic acids
O
(R C ONa)

Burning Burns with a non
luminous blue
A saturated
hydrocarbon with low
carbon content
C C
Burns with a luminous
yellow and sooty flame
An unsaturated
organic compound
C =C or C C

Heating a
mixture of the
test substance
with sodalime
Colourless gas which is
evolved burns with a
luminous flame
A salt of a
carboxylic acid
R-C-ONa
The gas evolved is
an alkane
(saturated)
R-C- or C-

Adding sodium
carbonate to
the test
substance
Colourless gas is
evolved which forms a
white precipitate with
lime water
Compound is a
carboxylic acid or
ester which is
hydrolysed
O O
R-C-OH or R- C-O-R
No reaction occurs Compound not a
carboxylic acid or
an ester

Adding acidified
potassium
manganate (VII)
Decolourisation occurs
in the cold
Unsaturated
compound or an
alkanol ( areducing
agent) c=c , c c
or R OH
Decolourisation occurs
on warming
Oxalates
O O
-O – C- C - O- or
oxalic acid
H2C2O4
Decolurisation occurs Unsaturated
compound
C =C or –C=C-

Adding
carboxylic acid
followed by
sulphuric (VI)
acid then
warming the
mixture
Oily pleasant smelling
compound formed
The compound is
an alkanol.
R-OH

Note:
•Potassium chromate (VI) could be used in all cases where
potassium manganate (VII) is used. In all these cases, orange
chromate (VI) ion changes to green chromate (III) ion.
•Tests on organic compounds are based on the physical of the
compounds: that is colour, smell, physical state at room
temperature, and the solubility in water. The tests do not,
therefore, give much indication regarding the identity of the
compound being investigated; they give the functional
regarding the identity of the compound being investigated;
they give the functional groups series which is not a clear
answer in most cases.

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