Day 1. Solutions. Emulsions and colloids.

EXPERIMENT 1
Preparation of 0.2 M KCl solution

Molarity is one of the ways to express solution concentration.

Molarity is defined as the number of moles of solute per liter of solution.

Given: Find:
C = 0.2 M m=?
V = 100 mL = 0.1 L M=?
C = n/V C = m/MV
n = m/M m = CMV
M(KCl) = 39.1 g/mol + 35.45 g/mol = 74.55 g/mol
m = 0.2 mol/L · 74.55 g/mol · 0.1 L
m = 1.491 g

To make 0.1 L of a 0.2 M KCl solution, we add 0.02 mol of KCl to a flask and then
add water to make 0.1 L of solution

EXPERIMENT 2
Dilution of 1 M CH3COOH solution

When diluting acids, always add the concentrated acid to the water.
Never add water to concentrated acid solutions.

A stock solution is a concentrated solution used to prepare more dilute solutions.

The dilution equation:

C1V1 = C2V2
where C1 and V1 are the molarity and volume of the initial concentrated solution
C2 and V2 are the molarity and volume of the final diluted solution.
Given: Find:
C2 = 0.1 M C1 = 1 M
V2 = 100 mL = 0.1 L V1 = ?
C1V1 = C2V2
V1 = C2V2 / C1
V1 = 0.1 mol/L · 0.1 L / 1 mol/L
V1 = 0.01 L

EXPERIMENT 3
Preparation of emulsions (oil in water)

Water + oil or grease

The combination of water and oil did not mix. No matter how hard the mixture was
shook, the oil always separated when the mixture was allowed to stay. In other
words, the oil did not dissolve in water.

Two phases formed. The upper phase was the oil. The lower phase was the water.

Water with oil or grease + soap (emulsifier)

Water and soap mixed quite easily. When soap was added to a mixture of water and
oil, the oil was no longer separated out from the water. It broke up into tiny particles
that remained suspended in the water without dissolving in it.

The water, with the oil particles in it, constitutes an emulsion.

An emulsion is a liquid, containing small, undissolved fatty particles.

The role of soap as emulsifier!

The soaps are sodium or potassium salts of higher fatty acid.
A soap molecule is compound of hydrophobic (water repelling) and hydrophilic (water
loving) parts. At oil water interface the soap molecule orients in such a way that the
hydrophobic (non-polar) part is directed towards oil phase while the hydrophilic
(polar) –COO–Na+ is directed towards water phase. Many molecules of emulsifier
concentrate at the oil-water interface and form a film around oil droplets. This
aggregate is called a micelle. This film prevents the merging of droplets.

Micelle

EXPERIMENT 4
Preparation of colloidal sulfur.

Na2S2O3 + H2SO4 → Na2SO4 + H2S2O3

H2S2O3 → H2O + SO2 + S
The color of the formed colloidal sulfur is yellow.
Heating is a method of bringing about the coagulation of colloidal dispersions.
EXPERIMENT 5
Preparation of colloidal solution of iron(III) hydroxide (ferric hydroxide).

FeCl3 + 3H2O → Fe(OH)3 + 3HCl

The color of the formed sol - iron(III) hydroxide - is red-brown.

EXPERIMENT 6
Coagulation of colloidal solution of iron(III) hydroxide.

Fe(OH)3 + NaCl → x
Fe(OH)3 + Na2SO4 → x

Sol - iron(III) hydroxide - can be made to aggregate by the addition of an ionic

solution, particularly if the solution contains anions with multiple charges.
Colloidal particles of the iron(III) hydroxide are positively charged, so the greater of
the negative charge, the more effective is the coagulation.
 DAY 2. Separation of Mixtures

EXPERIMENT 7
Soxhlet Extraction

Student should: - give the brief description of the procedure.

- sketch Soxhlet apparatus and label it correctly

EXPERIMENT 8
Fractional distillation

Student should: - give the brief description of the procedure.

- sketch distillation apparatus and label it correctly

EXPERIMENT 9
Separation of a mixture of copper(II) sulfate and benzil.

Student should: - give the brief description of the procedure.

- sketch separator funnel and label phases correctly

Obs.:
upper phase - copper(II) sulfate dissolved in water (blue colour)
lower phase - benzil dissolved in dichloromethane (yellowish green)

EXPERIMENT 10
Filtration

CaCl2 + K2C2O4 → CaC2O4↓ + 2KCl

Obs.: The color of the formed precipitate – calcium oxalate - is white.

Student should: - give the brief description of the procedure.

- sketch (gravity, vacuum) filtration apparatus and label them correctly
 DAY 3. The preparation and properties of selected chemical elements

EXPERIMENT 11
Laboratory preparation of hydrogen.
Reaction of zinc with hydrochloric acid.

Aim: to prepare hydrogen in lab and study its properties.

Reaction:
Zn + 2HCl → ZnCl2 + H2↑

Observation:
Bubbles of H2 hydrogen gas are seen evolving briskly. The burning splinter (brought
near the mouth of the test-tube) gets extinguished and the collected gas burns with a
popping sound. The produced H2 hydrogen gas is colourless, odourless and
tasteless.

Physical properties of hydrogen:

- it does not dissolve in water and does not support combustion.
- it is a highly inflammable gas
- it burns with a blue flame and produces a pop sound
- it is lighter than air
- it is neutral to litmus
------------------
- it is used as a reducing agent.

EXPERIMENT 12
Laboratory preparation of oxygen.
Thermal decomposition of potassium manganate.

Aim: to prepare oxygen in lab and study its properties.

When potassium permanganate is heated, it decomposes to give potassium

manganate, manganese dioxide and oxygen.

Reaction:
2KMnO4 → K2MnO4 + MnO2 + O2↑

ox. state of Mn in KMnO4: +7 (violet)

ox. state of Mn in K2MnO4: +6 (green)
ox. state of Mn in MnO2: +4 (brown)
ox. state of Mn in MnSO4: +2 (pink, colorless)

Observation:
The produced O2 oxygen gas is colourless, odourless and tasteless.

Physical properties of oxygen:

- it is slightly soluble in water
- it is not combustible but it supports combustion
- it is slightly heavier than air
- it is neutral to litmus
------------------
- it is used as an oxidizing agent

Tests for oxygen:

Oxygen gives reddish brown vapours with nitric oxide gas.
Oxygen turns hot pyrogallol solution to brown colour.

EXPERIMENT 13
Burning elements in oxygen. Combustion reaction.

Aim: to study properties of oxygen.

When elements burn in oxygen, they form oxides.

Metals form solid oxides.

They give alkaline solutions if they dissolve in water.
They turn red litmus blue.

Non-metals form oxides which are solid, liquid or gaseous.

These oxides are acids.
They turn blue litmus red.

a)
The burning splinter (brought near the mouth of the test-tube) continues to burn very
brightly in O2 oxygen gas and a glowing splinter is rekindled.
C + O2 → CO2 carbon dioxide, an invisible gas
It glows red.
b)
S + O2 → SO2 sulfur dioxide, a fuming gas with a choking smell
It burns with a blue flame.

c)
2Mg + O2 → 2MgO magnesium oxide, a white solid
It burns with a bright white flame.

d)
2Cu + O2 → 2CuO copper(II) oxide, a black solid
It does not burn. It turns black.

EXPERIMENT 14
Preparation of metallic copper.

Aim: to prepare metallic copper in lab.

to demonstrate that a more reactive metal can replace a less reactive metal
from its solution (the reactivity series of metals)
If you place a clean iron nail into a beaker of copper sulfate there is an interesting
reaction. The blue copper sulfate solution changes to a slightly paler colour. The
most remarkable thing that happens is that the nail looks a different colour. It has
become a copper colour.

CuSO4 + Fe → FeSO4 + Cu

The iron nail has become coated with copper. Iron is more reactive than copper and it
has pushed out copper from the copper sulfate and has reacted to form iron sulfate.
This pushing out is called displacement, so this type of reactions is named a
displacement reaction. A more reactive metal can replace a less reactive one in a
salt.

The solution of copper sulfate reacts with the iron nail to produce metallic copper.

EXPERIMENT 15
Flame test analysis.

Li+ - red
Na+ - yellow
K+ - pale lilac
Ca2+ - brick red
Sr2+ - crimson red
Ba2+ - apple-green
Cu2+ - blue green

Cobalt glass is used in flame test analysis.

It - absorbs the yellow light of sodium
- allows the much less intense lilac colour of potassium to be seen.
 Day 4. Acids, bases and salts – preparation, properties, types of
oxides
EXPERIMENT 16
Preparation of chromium(III) hydroxide. Amphoteric properties of Cr(OH)3.

Aim: to prepare chromium(III) hydroxide and study its amphoteric properties

Cr2(SO4)3 + 6NaOH → 2Cr(OH)3↓ + 3Na2SO4

Obs.: gelatinous grey-green precipitate of Cr(OH)3 is formed

Cr(OH)3↓ + NaOH → Na[Cr(OH)4]

Cr(OH)3↓ + 3NaOH → Na3[Cr(OH)6]
Obs.: precipitate of Cr(OH)3 is soluble in excess of NaOH
giving green solution

Cr(OH)3↓ + 3HCl → CrCl3 + 3H2O

Obs.: precipitate of Cr(OH)3 is soluble in HCl
giving green solution

2Cr(OH)3↓ + 4NaOH + 3H2O2 → 2Na2CrO4 + 8H2O

Obs.: Cr(OH)3 is oxidized by H2O2 and NaOH to Na2CrO4 and dissolves
giving yellow solution

EXPERIMENT 17
Preparation of silver salts

Aim: to prepare silver salts and study their properties

AgNO3 + KCl → AgCl↓ + KNO3

Obs.: white precipitate of AgCl is formed

AgNO3 + KBr → AgBr↓ + KNO3

Obs.: very pale cream precipitate of AgBr is formed

AgNO3 + KI → AgI↓ + KNO3

Obs.: very pale yellow precipitate of AgI is formed

Obs.: all of the precipitates change color if they are exposed to light
- taking on grey or purplish tints.

AgCl↓ + 2NH3·H2O → [Ag(NH3)2]Cl + 2H2O

Obs.: white precipitate of AgCl dissolves in ammonia solution
to give a colorless solution

[Ag(NH3)2]Cl + 2HNO3 → AgCl↓ + 2NH4NO3

Obs.: white precipitate of AgCl forms
EXPERIMENT 18
Reactions of metal oxides with water

Aim: to study the reactivity of metal oxides towards water and their properties

Phenolphthalein Methyl orange

CaO + H2O → Ca(OH)2 turns pink yellow
Obs.: basic oxide
CrO3 + H2O → H2CrO4 colourless pink
Obs.: acidic oxide
FeO + H2O → x x x
ZnO + H2O → x x x

EXPERIMENT 19
Reactions of metals with bases

Aim: to study the reactivity of metals towards sodium hydroxide

1) Mg + NaOH → no reaction

2) Zn + 2NaOH + 2H2O → Na2[Zn(OH)4] + H2↑

granular zinc dissolves in NaOH
to form tetrahydroxozincate, [Zn(OH)4]2–, ions

EXPERIMENT 20
Reactions of metals with acids

Aim: to study the reactivity of metals towards acids

with diluted H2SO4

Cu + H2SO4 → no reaction
Obs.: copper does not react with dilute sulfuric acid

a) cold H2SO4 Fe + H2SO4 → FeSO4 + H2↑

Obs.: iron reacts with dilute sulfuric acid to form iron(II) sulfate
(colorless solution) and to liberate hydrogen
the hydrogen formed burns with a pop sound
when a lighted matchstick is brought near it

with concentrated H2SO4

Cu + 2H2SO4 → CuSO4 + SO2↑ + 2H2O

Obs.: copper reacts with concentrated sulfuric acid to form copper(II) sulfate
(blue solution) and to liberate sulfur dioxide

a) cold H2SO4 Fe + H2SO4 → no reaction

Obs.: iron does not react with cold concentrated sulfuric acid

b) hot H2SO4 2Fe + 6H2SO4 → Fe2(SO4)3 + 3SO2↑ + 6H2O

Obs.: iron reacts with hot concentrated sulfuric acid to form iron(III) sulfate
and to liberate sulfur dioxide

with concentrated HCl

Cu + HCl → no reaction
Obs.: copper does not react with concentrated hydrochloric acid

with 2M HNO3
3Cu + 8HNO3 → 3Cu(NO3)2 + 2NO↑ + 4H2O
Obs.: copper reacts with 2M nitric acid to form copper(II) nitrate
(blue solution) and to liberate nitrogen oxide

with concentrated HNO3

Cu + 4HNO3 → Cu(NO3)2 + 2NO2↑ + 2H2O
Obs.: copper reacts with concentrated nitric acid to form copper(II) nitrate
(blue solution) and to liberate nitrogen dioxide (brown gas)

EXPERIMENT 21
Reactions of selected salts with acids

Aim: to study the reactivity of salts towards acids

Na2SO4 + CH3COOH → no reaction

2CH3COONa + H2SO4 → Na2SO4 + 2CH3COOH

Na2SO4 + HCl → no reaction

Na2CO3 + H2SO4 → Na2SO4 + H2CO3

H2CO3 → H2O + CO2↑

NaCl + H2SO4 → NaHSO4 + HCl

2NaCl + H2SO4 → Na2SO4 + 2HCl

NaNO3 + H2SO4 → NaHSO4 + HNO3

2NaNO3 + H2SO4 → Na2SO4 + 2HNO3

Na2SO4 + HNO3 → no reaction

 DAY 5. Types of chemical reactions

EXPERIMENT 22
Oxidation of copper metal with concentrated nitric acid to copper(II) nitrate

Cu + 4HNO3 → Cu(NO3)2 + 2NO2↑ + 2H2O

Obs.: copper reacts with concentrated nitric acid to form copper(II) nitrate
(blue solution) and to liberate nitrogen dioxide (brown gas)

EXPERIMENT 23
Precipitation of copper(II) hydroxide with sodium hydroxide

Cu(NO3)2 + 2NaOH → Cu(OH)2↓ + 2NaNO3

Obs.: copper reacts with sodium hydroxide to form
the blue gelatinous precipitate of copper(II) hydroxide

EXPERIMENT 24
Conversion of copper(II) hydroxide to copper(II) oxide

Cu(OH)2↓ → CuO + H2O

Obs.: the blue gelatinous precipitate of copper(II) hydroxide
decomposes to black precipitate of copper(II) oxide

EXPERIMENT 25
Dissolving of copper(II) oxide with sulfuric acid

CuO + H2SO4 → CuSO4 + H2O

Obs.: the black precipitate of copper(II) oxide
dissolves in sulfuric acid to form the blue solution of copper(II) sulfate

EXPERIMENT 26
Reduction of Cu2+ ions with zinc metal

CuSO4 + Zn → ZnSO4 + Cu
Obs.: the blue solution of copper(II) sulfate reacts with the iron nail
to produce colorless solution of zinc(II) sulfate and metallic copper

EXPERIMENT 27
Oxidation of carbohydrates with Cu(OH)2 (Trommer's test)

CuSO4 + 2NaOH → Cu(OH)2↓ + Na2SO4

2Cu(OH)2 + glucose → 2CuOH + H2O + gluconic acid
2CuOH → Cu2O + H2O

Obs.: copper hydroxide(II) Cu(OH)2 is reduced by glucose

first to CuOH (orange precipitate), then to Cu2O (red precipitate)
glucose is oxidized by copper hydroxide(II) Cu(OH)2 to gluconic acid