‫ﻣﮭﺪاه ﻟﺮوح اﻟﻤﺮﺣﻮم اﻟﺪﻛﺘﻮر ﻋﻤﺮو ﺧﺎﻟﺪ ﻗﻨﺪﯾﻞ‬
Topic 1: The particulate nature of matter
• Solids, liquids and gases
• Diffusion
Topic 2: Experimental techniques and chemical analysis
•
•
•
•
Experimental design
Safety in laboratory
Separation and purification
Chromatography
Topic 3: Atoms, elements and compounds
•
•
•
•
•
•
Elements, compounds and mixtures
Atomic structure and periodic table
Isotopes
Ions and ionic bonds
Simple molecules and covalent bonds
Giant covalent structures
Topic 4: Stoichiometry
• Chemical formula
• Relative masses of atoms and molecules
• The mole and the Avogadro constant
Topic 5: The periodic table
•
•
•
•
•
•
Arrangement of elements
Group I properties (The alkali metals) properties
Group VII (The halogens) properties
Transition metals
Nobel gases
Redox
Topic 6: Electricity and chemistry
•
•
Electrolysis
Extraction of aluminium
1
Topic 7: Acids, bases and salts
• The characteristic properties of acids bases
• Oxides
• Preparation of salts
Topic 8: Metals
• Properties of metals
• Uses of metals
• Alloys and their properties
• Reactivity series
• Corrosion of metals
• Extraction of metals
Topic 9: Chemical energetics
• Exothermic and endothermic reactions
Topic 10: Chemical reactions
• Physical and chemical changes
• Rate of reaction
• Reversible reaction and equilibrium
Topic 11: Chemistry of the environment
• Water
• Fertilizers
• Air quality and climate
Topic 12: Organic chemistry
•
•
•
•
•
•
•
•
Formulae, functional groups and terminology
Naming organic compounds
Fuels
Alkanes
Alkenes
Alcohols
Carboxylic acids
Polymers
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Grade (10)
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Grade (10)
Topic 1
States of Matter
“Matter”
­ Is anything has mass and occupying a space.
­ Is made of tiny particles having energy which causes them to vibrate and / or move.
­ There are three states (phases) of matter: solid, liquid and gas.
Solid
Liquid
Gas
Lattice: regular 3D
arrangement of particles
in a crystalline solid.
Arrangement of ­ Closely packed
particles
­ Regular arrangement
of particles (lattice)
Intermolecular ­ Negligible
spaces
Intermolecular ­ Very strong
forces
Movement of
­ Vibrate in a fixed
particles
position
­ Touched
­ Randomly
distributed
­ Very small
­ Very far apart
­ Totally randomly
distributed
­ Very large
­ Weaker than in
solids
­ Slide over each
other
­ Very week
Shape
­ Fixed shape
­ Crystalline lattice
Volume
­ Fixed
­ No fixed shape
(takes the shape of
the container)
­ Fixed
Compression
­ Can not be
compressed
­ Can be hardly
compressed
­ Free to move
randomly in all
directions
­ No fixed shape
­ No fixed volume
(volume of the
container)
­ Can be
compressed
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­ The movement of the particles depends on:
1. Mass of particles.
2. Their kinetic energy.
“Changes of State”
Energy is given in
Evaporation
Solid
Liquid
Condensation
Gas
Energy is given out
N.B. Examples of materials which sublime:
- Iodine is a dark gray solid that sublimes to purple vapour
-White solid CO2 (dry ice) that sublimes to CO2 gas.
-Naphthalene (moth balls) & ammonium chloride.
Factors affecting evaporation:
1-Temperature: higher temperature leads to faster evaporation
2­Volatility: volatile liquids (having low b.p.) like alcohols evaporate faster
than nonvolatile liquids like water.
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3­Surface area: larger surface area leads to faster evaporation.
N.B. 1. When pressure increases over a liquid, the boiling point increases.
2. Evaporation occurs due to moving of air particles which possess large amount
of energy and hit the particles of a liquid giving them kinetic energy enough to overcome
the forces of attraction, escape and diffuse as a gas into the surrounding medium
­ Kinetic particle theory:
When a solid is heated, the particles vibrate faster about a fixed point. This causes the
particles to move further apart and so the solid expands.
When the particles gain sufficient energy to overcome the strong forces of attraction
holding them together, they can move out of their fixed positions. They can slip and
slide over each other in a continuous random motion. When this happens the solid
melts.
The particles in the liquid are still close to each other. They have enough kinetic energy
to move around each other closely, but do not have enough energy to overcome the
forces that hold them close to each other.
If more heat energy is supplied to the particles, they move faster until they have enough
energy to overcome the forces holding them together. The particles then escape from
the liquid surface and move around in a continuous rapid random motion. The liquid
now boils.
In the vapour formed, the particles move in a rapid random motion. The movement is
random due to the collision of the vapour particles with the air particles.
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Heating Curve
Melting point is the temperature at which solid changes to liquid.
Boiling point is the temperature at which liquid changes to gas.
Value of melting point and boiling point depends on the intermolecular forces
between particles.
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N.B.
During melting process the temperature remains constant as the energy is used to
overcome the forces holding the lattice.
During boiling process temperature remains constant as the energy is used to
overcome the forces between the particles.
Remember, evaporation occurs at room temperature, while boiling needs heating till
boiling point.
Any change in state is:
- a physical change such that no new substance is formed.
- a reversible change.
- does not affect the mass of the substance.
“Diffusion”
• Diffusion is the spreading of gas or liquid particles from more concentrated area to
less concentrated one.
• The particles mix and spread by colliding with other moving particles and bouncing
off in all directions.
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Rate of diffusion depends on:
1) Mass:
­ Less dense particles of lower atomic or molecular mass (Ar or Mr)] diffuse faster than
more dense particles of higher atomic or molecular mass at same temperature.
2) Energy:
­ Particles with more kinetic energy diffuse faster than particles of less K.E with the
same mass.
3) Presence of other substance:
­ Diffusion takes place faster in vacuum ( no other substances )
4) Intermolecular spaces:
­ Diffusion takes place faster in gases than in liquids ( gases have larger intermolecular
spaces )
N.B. No diffusion in solids.
(A) Diffusion in gases:
1. Diffusion of bromine:
The reddish brown bromine gas move randomly, collide with air particles diffuses
upwards between air particles, mix and spreads uniformly to fill both gas jars.
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2. Diffusion of ammonia and hydrogen chloride gas:
A + B
­ White cloud of NH4Cl is formed nearer to the cotton wool soaked with HCl [more
dense / more relative molecular mass].
­ The ring is not formed immediately because
1­ The particles are not moving in just one direction
2­ The tube is filled with air
NH3 molecules have less mass than the HCl molecule, so diffuse faster, hence the product (a white
cloud of NH4Cl) forms closer to the end where the HCl is
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(B) Diffusion in Liquids:
Diffusion of copper sulfate in water:
­ Blue crystals dissolve, sulfate particles fill inter molecular spaces of water.
­ Both water and sulfate particles are in a continuous random motion& collide.
­ Blue colour of copper sulfate spreads gradually as the blue particles diffuse in
water.
­ Water becomes uniformly blue.
The particles could be:
­ Atoms that cannot break down further in chemical reaction ex. He
­ Molecules that consist of two or more atoms joined together ex. Br2.
­ Ions that are charged atoms or groups of atoms ex. Ions in CuSO4
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Gas pressure:
Gas pressure is due to the collision of gaseous particles with the walls of the
container.
­ When the gas is heated in a closed container, the particles move faster hit the walls
more often and with more force, its pressure increases ex. pressure cooker.
­ When the gas is compressed into a smaller space, particles hit the walls more often,
its pressure increases.
­ When the gas is heated, particles gain more energy move faster and away from
each other, volume increases.
­ When the gas is cooled, particles loose more energy move slower and near from
each other, volume decreases.
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Topic 2
Experimental Techniques
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Grade (10)
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Grade (10)
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Grade (10)
“Measurements”
1. Volume:
a) Gas
is measured using gas syringe (cm3), or measuring cylinder.
b) Liquid:
* Accurate burette, volumetric pipette or graduated pipette (cm3)
* Rough measuring cylinder (cm3)
c) Solid
measuring cylinder [VSolid = V2 – V1]
2. Time:
Stop clock
places
seconds or minutes which is usually accurate to one or two decimal
3. Mass:
Top pan balance (digital balance)
which normally give readings to two decimal
places. These must be tarred (set to zero) before use gram
4. Temperature:
Thermometer
can normally give readings to the nearest degree Celsius ºC
“Apparatus in the Lab.”
• Test tube: used for any chemical reaction.
• Filter funnel: to filter solution.
• Burette : is the most accurate way of measuring a variable volume of liquid between
0 cm3 and 50 cm3 (e.g. in a titration)
• Volumetric pipette , usually 10 cm3 or 25 cm3( multiple of 5)
• Measuring cylinder: used to measure approximate volumes where accuracy isn´t an
important factor. These are graduated (have a scale so can be used to measure) and
are available in 25 cm3, 50 cm3, 100 cm3 and 250 cm3
• Test tube Rack: to store tubes.
• Beaker: used for dissolving.
• Bunsen burner, Tripod and Gauze: strong heating system.
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• Combustion spoon: used for combustion.
• Watch glass: for covering and reduces evaporation.
• Thermometer: used to measure temperature.
• Unglazed porcelain: Anti-bombing granules can be used to ensure smooth boiling of
liquid.
• Glass rod: For stirring and does not involve in the reaction.
Diagram of the set-up for an experiment involving gas collection
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Grade (10)
Diagram of the set-up for an experiment involving gas collection
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“Matter”
Mixture
Pure
Made of two or more substances mixed together
Made of one type of
(not bonded) in any ratio
particles
Element
Made of only one
type of atoms
Compound
Made of two or more
elements chemicallybonded together in
fixed ratio
Monoatomic
Diatomic
Ex. Inert gases
Ex. Cl2, H2, O2
N.B.
Impurity is the unwanted substance, mixed with the substance you need.
Medical drugs, water and food flavoring must not contain any impurities that could
harm people (must be safe).
To make sure of the purity of a certain substance, measure its m.p & b.p.
Pure substance has a definite, sharp m.p & b.p.
When the substance has impurities its melting point falls and its b.p rises and both
will be over a range.
The more impurity there is, the bigger the change in m.p and b.p, and the wider the
range over which melting and boiling occur.
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Differences between mixtures and compounds:
Mixture
1- It contains two or more different
substances in any ratio.
2- No chemical change takes place
Compound
- 1- It is a single substance made of
two or more different elements
chemically bonded in a definite
ratio.
2- Involves a chemical change.
3- The components can be separated
by physical means
3- The components can be separated
by chemical means.
4- Keeps the properties of their
components.
4- Its properties are different from
those of its components.
5- No change in energy when the
mixture is formed.
5- Energy is given out or absorbed
6- Ex. Fe / S
6- Ex. FeS
“Types of Mixtures”
(1) Solid / Liquid Mixture.
(2) Liquid / Liquid Mixture.
(3) Gas / Gas Mixture.
(4) Solid / Solid Mixture.
♣ How to separate the components of a mixture?
(1) Solid / Liquid Mixture
Soluble
1) Evaporation.
2) Simple distillation.
3) Crystallization
Insoluble
1) Decantation:
- Pouring the liquid off the insoluble substance
[big particles]
2) Filtration:
- Sand & Water [small particles]
3) Centrifugation:
- The sample is spun round very fast and the solid is
flung to the bottom of the tube [tiny particles (blood)]
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Solution (salty water) = Solute (table salt) + Solvent (water)
• Solution: is a mixture of solute and solvent.
• Solute: is the substance that dissolves in solvent to form solution.
• Solvent: is the substance that used to dissolve the solute [Ex. water , ethanol]
•
•
•
•
Solutions can be either:
Diluted solution: Small amount of solute / 1 dm3 solution.
Concentrated solution: Large amount of solute / 1 dm3 solution.
Saturated solution: Formed when no more solute can dissolve in the solution
at a certain temperature.
Solubility: is the maximum amount of solute in gram which dissolves in 100 g of
solvent at a given temperature.
Solubility curve:
These are curves that show how the solubility of a solid changes with temperature
At 20 oC the solubility of KNO3 is 32 g / 100 g water.
At 60 oC the solubility of KNO3 is 110 g / 100 g water.
Dissolving increases by:
1. Heating
2. Stirring
3. Crushing the solute [large surface area]
N.B. Solubility of gases in liquids decreases by heating
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1) Evaporation:
The solution is boiled till dryness to evaporate the solvent, the powder solid is left
behind.
2) Simple distillation:
- A way to obtain the solvent from a solution.
- The solution is heated till it boils, turns to vapour, and rises into the condenser.
- The solvent is condensed back to a pure liquid and collected, the salt is left behind.
3) Crystallization:
Heat till point of crystallization. Leave to cool, filter and dry between two filter paper.
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N.B.
Saturated solution must be left at room temperature to cool down ,to get large
crystals
Do not:
1. stir
2. put in refrigerator
1) Decantation: [big particles]
- Pouring the liquid off the insoluble solid.
2) Filtration:
[smaller particles]
• Residue is a substance that remains after filtration,
evaporation or distillation.
• Filtrate is a liquid r solution that has passed
through a filter
• Filter off the insoluble solid. Rinse with distilled water
to remove the soluble substance.
3) Centrifugation:
[very small particles]
- The sample is spun round very fast and the solid is flung to the bottom of the tube.
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(2) Liquid / Liquid Mixture
Miscible liquids
Immiscible liquids
- They do not mix easily.
- They are uniformally mixed.
- They are separated using separating
- Their
separation
depends
on
funnel according to the difference in
difference in boiling point. (Ex. water
their densities. (Ex. oil & water)
& ethanol)
[by fractional distillation]
Fractional Distillation
- Used to separate liquids with different b.p.
- The liquid with the least b.p distills first.
Uses of frictional distillation:
1. To separate liquids from each other.
2. To separate fractions of crude oil.
3. To separate gases such as nitrogen from liquid air, the gases boil off one by one.
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(3) Gas / Gas Mixture
• Air is a mixture of gases; its components can be separated by fractional distillation of
liquid air.
• Air is liquefied by applying high pressure and low temperature, then allowed to warm
up :
Nitrogen boils first at -196 o C.
Argon boils second at -186 o C.
Oxygen boils third at – 183o C.
• Diffusion as a less dens gas diffuses faster, ex. H2 and CO2
(4) Solid / Solid Mixture
• By magnet [magnetic property]:
Ex. (Fe , Co , Ni)
- Iron /sulfur mixture can be separated by magnet.
• Solvent extraction [solubility]:
Ex. (sand / table salt)
Steps:
* crush the mixture
* add water
* stir with gentle heating
* filter
((Filter))
Filtrate
Residue (sand)
- Heat the filtrate till crystallization point, leave to cool to get crystals or
- Evaporate the filtrate till dryness to get powder
- Mixture of salt and sugar can be separated by dissolving in ethanol (water dissolves
both). Sugar dissolves in ethanol but not salt. Ethanol is flammable, so should be
evaporated on water bath.
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• Chromatography:
- Used to separate a mixture of substances [in small amount].
N.B.
1. Draw the base line in pencil which does not produce spots (ink is not used as it
leaves spots).
2. Apply the spot on the base line, and put the paper in the solvent such that its level is 2
cm below the base line
3. Leave till the solvent reaches near the end of the paper, remove it and dry
4. If the spot is from one substance, it will leave one spot.
The use of chromatography :
1. Testing for the purity of substances.
2. Discover the substances present.
3. In medical labs.
The separation by chromatography is due to:
1. Different solubility of the components in the solvent.
2. Different degree of diffusion through the chromatogram paper (capillarity).
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• Number of spots = Number of components
• Proteins separate to amino acids and carbohydrates separate to glucose sugar, both
are colourless, so we spray the spots with locating agent to make them visible.
• To identify the spots ,we refer to control or calculate the flow rate Rf (retention
factor)
Y
X
Distance moved by the substance
Rf =
<1
Distance moved by the solvent
Rf
X
= ---Y
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Using Rf values to identify components of a mixture
N. B. To get a chromatogram of chlorophyll follow the following steps:
1. Crush small pieces of green leaves with sand to increase friction and get more extract
2. Add ethanol to dissolve chlorophyll then filter
3. Concentrate the solution over water bath as ethanol is flammable
4. Run chromatography you will get two spots
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Topic 3
Atoms, Elements and compounds
“The Atom”
­ The atom is the smallest building unit of an element that take part in chemical
reactions and cannot be splitted into anything simpler.
• The atom consists of:
1. Nucleus at the center (massive)
[p+ , n0] ∴ The nucleus is positively charged.
2. Negative electrons rotate around the nucleus in energy levels (shells & orbits).
[p+ , n0 , e–] are subatomic particles.
×
● p+
○ n0
× e–
×
The mass of the atom is concentrated in the nucleus as mass of (e–) is negligible if
compared to the mass of (p+) and (n0).
Particle
Proton
Neutron
Electron
Symbol of element
Nucleon number = number of p+ + n0
(Mass number)
Example: 7Li3
n0 = 7 – 3 = 4
Symbol
p
n
e
Mass (amu)
1
1
1/1840
X
A
Z
Charge
+1
0
–1
Proton number = number of p+
(Atomic number)
[proton number = 3 , nucleon number = 7]
(p+ = 3 , e– = 3) ∴
N.B. (number of p+ = number of e–), so the atom is electrically neutral.
Nucleon number is the sum of Number of protons and neutrons inside the nucleus of
one atom.
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Proton number is the Number of protons inside the nucleus of one atom
The heaviest atom has 7 energy levels or electron sells which have different energies.
Energy of electrons increases as we go far from the neucleus.
K L M N O PQ
+
1 2 3 4 5 6 7
2e– 8e– 18e– 32e–
The last energy level cannot hold more than 8 electrons.
Na11 (2, 8, 1)
×
××
×
× ×
× ×
+
× ×
× ×
×
××
××
17
Cl (2, 8, 7)
×
××
××
×
+
×
××
××
××
××
××
×
+
×
××
××
×
×
× ×
× ×
×
××
×
+
×
××
×
×
× ×
× ×
K19 (2, 8, 8, 1)
Fe26: 2, 8, …, 2
As33: 2, 8, 18, … Sr38: 2, 8, …, 8, 2
× ×
× ×
× ×
× ×
Ca20 (2, 8, 8, 2)
Pd46: 2, 8, 18, …, 2
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“Isotopes”
1H
1
p+ = 1
e– = 1
n0 = 1 – 1 = 0
(Protium)
2H
1
3H
p+ = 1
e– = 1
n0 = 2 – 1 = 1
(Deuterium)
17
35Cl
1
p+ = 1
e– = 1
n0 = 3 – 1 = 2
(Tritium)
17
37Cl
• Isotopes: are atoms of same element having same number of protons [p+] but
different number of neutrons [n0].
N.B.
Isotopes of an element have
1-Same chemical properties (as they have same Number of electrons. Same
electronic configuration and same valence electrons.)
2-Different physical properties e.g. density, rate of diffusion (different atomic mass)
Part of the definition of relative atomic mass is ‘the average mass of naturally occurring
atoms of an element. Some relative atomic masses are not whole numbers.
How to calculate the relative atomic mass of an element
1. Element Y has only two different types of atoms present in an element Y is shown
35
Y: 37Y = 3:1. Calculate the relative atomic mass of element Y to one decimal place.
The answer:
2. The element Gallium has two isotopes 69Ga and 71Ga with abundance 60% and 40%
respectively. Calculate Ar of Gallium.
The answer:
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6
7
3. Lithium has two isotopes, 3Li and 3Li . The relative
abundance of these two isotopes is shown in the figure.
Calculate the relative atomic mass of lithium.
The answer:
4. The isotopes of magnesium and heir abundance are given in this table
Isotope
Symbol
Abundance / %
Magnesium ­ 24
24 / 12 Mg
78.6
Magnesium ­ 25
25 / 12 Mg
10.1
Magnesium ­ 26
26 / 12 Mg
Calculate the relative atomic mass of magnesium.
The answer:
11.3
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333333.
Grade (10)
The Periodic Table”
The elements are arranged according to the increase of their proton number
­ The group number indicates the number of electrons in the last energy level
(valence e–) in groups from I tp VII.
­ The period number indicates number of energy level occupied in one atom .
­ Group VIII nobel gases have full outer shell
Group
I
II
III
IV
V
VI
VII
0
Valence electrons
1
2
3
4
5
6
7
0
Valency
1
2
3
4
3
2
1
0
• Group number (valence electrons) = Valency = Number of electrons lost
Groups I,
II and III.
• Group number (valence electrons)
Valency = Group number ­ 8 = number of electrons gained or shared
Groups IV to
VII
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Chemical Bonding
• Except noble gases, elements react with each other:
In order to reach stability, to obtain outer most energy level filled with electrons and
have the configuration of nearest noble gases.
“Elements”
Metals
Non-metals
[1, 2, 3 electrons in outer most energy level]
[4, 5, 6 or 7 electrons in outer most energy level]
Lose electrons
Gain electrons
Share electrons with
(Form +ve ions)
(Form –ve ions)
non­metal or H atom
Electrostatic attraction force
1 pair
single / 2 pairs
double / 3 pairs
triple
(oppositely­charged ions)
[ionic bond]
[covalent bond]
“Ionic Bond”
­ Is a strong electrostatic attraction force between +ve ion (cations) and –ve ions
(anions) due to transfer of electrons from metal to non­metal.
Example:
(NaCl)
(MgO)
(CaCl2)
Na11: 2, 8, 1 & Cl17: 2, 8, 7
Mg12: 2, 8, 2 & O8: 2, 6 Ca20: 2, 8, 8, 2 &Cl20: 2, 8 7
××
Na+
●
×
Cl ××
–
××
Mg2+
●
●
××
O ××
××
2–
××
Ca2+
–
2 ●× Cl ××
××
• Ion: is a charged atom or group of atoms formed by the gain or loss of electrons
(unequal number of protons and electrons).
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Grade (10)
Ionic compound: a compound formed of oppositely charged ions joined by strong electrostatic
attraction forces.
•
Ionic structure [ionic lattice]: is a regular arrangement of oppositely charged ions, held together
by strong electrostatic attraction force (no molecules just ions).
• Number of positive charges on the positive ion(s) must equal number of negative
charges on negative ion(s), so the net charge of any ionic compound is zero ( no
molecules) .
Compound ions (atomic groups / radicals):
Compound ions is a group of different atoms chemically bonded carries positive or
negative charge and behave as one atom during chemical reactions
“Covalent Bond”
­ Is formed when atoms of non­metals share one or more pair of electrons forming
molecules.
N.B. No ions.
“Simple molecules”
Single covalent: Each atom shares by one electron [The bond is one pair of
electrons]
Example:
(H2)
(Cl2)
(HCl)
1H: 1
17Cl: 2, 8, 7
×
×
17Cl: 2, 8, 7
●●
××
Cl ●× Cl ●
●
H ×● Cl ××
××
H ×● H
1H: 1
××
●●
××
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(H2O)
1H: 1
8O: 2, 6
●●
×
H●
(CH4)
6C: 2, 4
1H: 1
●●
O
×●
H
Double covalent: Each atom shares by two electron [The bond is two pairs of
electrons]
(O2)
(CO2)
8O: 2, 6
××
×
×
6C: 2, 4
××
●●
O ●● ×× O
8O: 2, 6
×
×
●
●
××
O ●● ×× C ×× ●● O
×
×
(C2H4 )
Triple covalent: Each atom shares by three electron [The bond is three pairs of
electrons]
(N2)
7N: 2, 5
×
×
●×
●×
N ●× N
(C2H2)
6C: 2, 4
●
●
1H: 1
●●
H●× C ●● ●● C × ●H
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Differences between ionic and covalent compounds
Property
Elements in compound
Ionic compounds
Covalent compounds
Metal and non­metal
Two or more non­metals
Only ions
Uncharged molecules
Type of particle
(small or giant)
Volatility, melting point and
High melting and boiling points
Low melting and boiling points
boiling point
because ions are held together
because intermolecular forces
by strong electrostatic forces
are weak
Most are soluble in water but
Most are insoluble in water but
insoluble in organic solvents
soluble in organic solvents
Conduct when molten or
In general do not conduct as
dissolved in water because the
solids, liquids or in solution,
ions are free to move. Do not
because there are no ions, only
conduct as solids, because the
molecules. A few dissolve in
ions are stuck in the crystal
water and form ions, and these
lattice
will conduct in aqueous solution,
Solubility
Electrical conductivity
for example hydrogen chloride.
“Macromolecules”
Contain big number of atoms joined together by:
­ Covalent bond
(giant covalent structure)
• Allotropes: different structural forms of an element in the same state.
N.B. Diamond and Graphite are allotropes.
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IGCSE
Grade (10)
Giant covalent structures
Diamond:
­ Is a crystalline form of Carbon (C), each C atom is strongly bonded to four carbon
atoms by covalent bonds in a tetrahedral structure.
­
Tetrahedral Structure
Properties of diamond
Uses
Very high melting point
Hardest substance known
Drilling; cutting glass and metals
Does not conduct electricity
No free electrons
Colourless crystals that glitter jewellery
Graphite:
­ Each C atom is strongly bonded to 3 C atoms by covalent bond forming layers of
hexagons held by weak force of attraction.
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IGCSE
Grade (10)
roperties of graphite
Uses
Black shiny solid
Soft with a slippery almost soapy feel: the layers can slip over each ­As a lubricant
other because of the weak bonds between layers
­In pencils (mixture
of graphite and clay)
Good conductor of electricity because the electrons between the
layers are mobile
To make electrodes
High melting point because the strong bonds in the layers have to
break before the graphite can melt
Silicon IV oxide (silica or sand):
­ Each silicon atom is strongly bonded to four oxygen atoms and each oxygen atom is
strongly bonded to two silicon atoms in a tetrahedral structure.
O
Si
O O
O
­ Silicon dioxide has same chemical properties of diamond
­ Used in
1. Sandpapers, as it is hard and can scratch things.
i. bricks for lining furnaces, as it has high m.p.
ii. making glass and lenses, as it is hard and let light through.
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IGCSE
Grade (10)
Metallic bonding
It is the electrostatic attraction force between regularly arranged positive metal ions, and
a mobile sea of delocalized electrons.
Properties of metals:
1. Have high m.p and b.p. due to the strong attraction force between positive metal
ions and the freely moving electrons.
2. Conduct electricity due to the free moving electrons within the structure.
3. Good conductors of heat, as the freely moving electrons within the structure transfer
heat energy along the metal .
4. Malleable and ductile. If a force is applied, the metal rows of ions slide over each
other.
5. Have high density, as the ions are very closely packed.
40
IGCSE
Grade (10)
Topic 4
Stoichiometry I
is the ratio of the reactants and products in a balanced symbolic equation
“Chemical Formula”
­ Chemical formula is a set of chemical symbols to present a chemical substance.
­ Ex. (NaCl , H2O)
♣ How to write a chemical formula?
Magnesium chloride
Mg
Valency
2
Calcium oxide
Cl
Ca
1
MgCl2
2
O
2
CaO
N.B. Simplify.
Valency is the number of electrons gained or shared by one atom during a chemical
reaction.
Simple rules for naming compounds:
1) Compounds of 2 elements:
• Metal and non-metal:
­ The metal is put first and the ending of the non­metal changes to (ide), ex.
magnesium chloride.
• 2 non-metals:
­ If contains hydrogen it comes first, Ex. hydrogen sulfide, except ammonia NH3
­ If 2 non­metals, the element in the lower group comes first, ex. nitrogen dioxide.
­ If in the same group, the lower one comes first, ex. sulfur trioxide
• Common names:
­ Water, Ammonia.
N.B. Radical is a group of different atoms, linked together, carries +ve or -ve
charge, behaves as one atom during chemical reactions.
2) Compounds containing radicals with enough oxygen:
­ (OH–) hydroxide, (NO3–) nitrate, (SO42–) sulfate, (CO32–) carbonate, ( HCO3–)
hydrogen carbonate (PO43–) phosphate, (S2O32­)Thiosulfate
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IGCSE
Grade (10)
3) Compounds containing radicals with less oxygen:
­ (NO2–) nitrite, (SO32–) sulfite
Remember (NH4+) ammonium is a positive radical, without oxygen
Aluminum hydroxide
Al
Sodium sulfate
OH
Na
3
1
Al(OH)3
SO4
1
2
Na2SO4
• Acids: HCl , H2SO4 , HNO3
• Molecular formula: represents number and type of atoms in one molecule ex. C2 H6
• Empirical formula: shows the simplest whole number ratio of the different atoms
present in a substance ex. CH3.
♣ Write the chemical formula for:
1. Phosphoric acid
3. Iron (III) hydroxide
5. Manganese (IV) oxide
7. Copper (II) nitrite
2. Hydro bromic acid
4. Zinc (II) chloride
6. Potassium sulfate
8. Calcium nitride
“Equations for Chemical Reactions”
• A chemical equation: shows the reactants, products and their ratio involved in a
chemical reaction.
* Word equation
* Balanced symbolic
Ex. Sodium burns in chlorine to form sodium chloride
­ The word equation:
Sodium + Chlorine
­ Balanced symbolic equation: 2Na + Cl2
Sodium Chloride
2 NaCl
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IGCSE
Grade (10)
♣ Write the chemical equations if:
1. Aluminum reacts with oxygen to form aluminum oxide.
2. Calcium hydroxide reacts with nitric acid to form calcium nitrate and water.
N.B. State symbols can be written under the formula:
(s)
solid
(l)
liquid
(g)
gas
(aq) dissolved in water “aqueous”
2Mg(s) + O2(g)
2MgO(s)
• Relative atomic mass (Ar) of an element: is the average mass of the element’s
isotopes relevant to an atom of Carbon 12.
• Relative molecular mass (Mr): is the sum of the relative atomic masses of all atoms
present in one molecule.
One mole of the substance has a mass equal to the relative formula mass in
gram.
­ Ar of Carbon = 12
* mass of 1 mole of C = 12 g
­ Mr of Oxygen O2 = 2 × 16 = 32
* mass of 1 mole of O2 = 2 × 16 = 32 g
­ Mr of MgCl2 = 24 + 2 × 35.5= 95
* mass of 1 mole of MgCl2 = 24 + 2 × 35.5 = 95 g
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IGCSE
Grade (10)
Stoichiometry II
Mole: is the amount of substance that contains 6x1023 particles.
Mole – Mass:
Mass
Number of moles (n) = ------------------------------------Molar mass [Ar or Mr]
1.
Calculate the number of moles for the following :
a)
2.4 grams of magnesium.
……………………………………………………………………………………………
b)
13 grams of Zinc.
……………………………………………………………………………………………
c)
4.4 grams of carbon dioxide.
……………………………………………………………………………………………
d)
40 grams of sodium hydroxide.
……………………………………………………………………………………………
2.
Calculate the mass of the following:
a)
0.5 mole of NaOH .
……………………………………………………………………………………………
b)
0.1 mole of CuSO4.5 H2O.
……………………………………………………………………………………………
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IGCSE
c)
Grade (10)
0.25 mole of aluminium hydroxide.
……………………………………………………………………………………………
d)
0.4 moles of Fe SO4.7H2O.
……………………………………………………………………………………………
Mole - Number of particles:
Each one mole of any substance contains 6.02x1023 particles (Avogadro’s number)
Number of particles
Number of moles (n) =------------------------------------6.02 x1023
3.
Calculate the number of particles for the following:
a)
0.2 mole of magnesium.
……………………………………………………………………………………………
b)
O.4 mole of Mg(NO3)2.
……………………………………………………………………………………………
4.
For 12.5 grams of Calcium Carbonate, find :
a)
Number of moles.
……………………………………………………………………………………………
b)
Number of ions.
……………………………………………………………………………………………
45
IGCSE
Grade (10)
5.
For 4 grams of NaCl, find:
a)
Number of moles.
……………………………………………………………………………………………
b)
Number of ions.
……………………………………………………………………………………………
Mole – Volume:
Volume
Number of moles (n) =------------------------------------24
N.B. Measuring unit of volume (V) is dm3
X1000
[dm3
cm3]
/1000
Remember:
1 mole of hydrogen gas H2
Contains 6.02x1023 molecules
Occupies volume of 24 dm3
has molar mass of 2x1 = 2 g
6.
Calculate the number of moles for the following:
a)
2.4 dm3 of CO2.
……………………………………………………………………………………………
b)
480 cm3 of ammonia gas.
……………………………………………………………………………………………
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IGCSE
c)
Grade (10)
3
600 cm of Oxygen.
……………………………………………………………………………………………
d)
960 cm3 of SO3 gas.
……………………………………………………………………………………………
7.
Calculate the volume of:
a)
0.3 moles of carbon monoxide.
…………………………………………………………………………………………
b)
4.4 grams of nitrogen dioxide.
……………………………………………………………………………………………
……………………………………………………………………………………………
c)
20 grams of sulfur trioxide.
……………………………………………………………………………………………
……………………………………………………………………………………………
d)
8.8 grams of Hydrogen chloride gas.
……………………………………………………………………………………………
……………………………………………………………………………………………
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IGCSE
Grade (10)
Mole – concentration:
Concentration of a solution (c): is the amount of solute in moles or gram,
dissolved in 1 dm3 of the solution
Amount of solute [moles]
C [mol/dm ] = ---------------------------------------Volume of solution [dm3]
3
n=CxV
C [g/dm3]
=
C [mol/dm3] X Mr
8.
Calculate the concentration in mol/dm3 and g/dm3 for the following:
a)
80 g of sodium hydroxide dissolved in 2 dm3.
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
b)
98 g of sulfuric acid dissolved in 100 cm3.
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
c)
1000 cm3 of lithium chloride solution containing 2 moles of the salt.
……………………………………………………………………………………………
…………………………………………………………………………………………..
……………………………………………………………………………………………
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9.
How many moles are there in each of the following:
a)
300 cm3 of sodium sulfate from 2 mol/dm3 solution.
Grade (10)
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
b)
500 cm3 of ammonia from its 1 mol/dm3 solution.
……………………………………………………………………………………………
…………………………………………………………………………………………….
……………………………………………………………………………………………
c)
2500 cm3 of nitric acid from its 0.3 mol/dm3 solution.
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
d)
1000 cm3 of sodium fluoride from its 0.5 mol/dm3 solution.
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
10. How many grams of :
Sodium hydroxide is present in 100 cm3 of its 0.1 mol/dm3 solution.
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
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IGCSE
Grade (10)
Mole and chemical equations:
11. 200 g of pure calcium carbonate are heated strongly. Calculate:
The mass of calcium oxide and volume of carbon dioxide produced.
CaCO3
CaO +
CO2
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
12. Calculate the volume of 1 mol/dm3 solution of H2SO4 required to react with
6g of Mg, and the volume of H2 produced at r.t.p.
Mg + H2SO4
MgSO4 + H2
……………………………………………………………………………………………
…………………………………………………………………………………………….
……………………………………………………………………………………………
……………………………………………………………………………………………
…………………………………………………………………………………………….
……………………………………………………………………………………………
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IGCSE
Grade (10)
N.B. For reactions involving only gases, the volume ratio equals the mole
ratio in the balanced equation.
13. Calculate the volume of methane needed to react with 70 dm3 of oxygen.
CH4 + 2 O2
CO2 + 2 H2O
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
a)
The volume of carbon dioxide produced at r.t.p.
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
…………………………………………………………………………………………….
14. Find the volume of oxygen needed to react completely with 15 g of C2H6.
C2H6 + 7/2 O2
2 CO2
+
3H2O
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
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IGCSE
15.
Grade (10)
What volume of carbon dioxide at r.t.p. will be formed when 50 g of calcium
carbonate react with an excess of hydrochloric acid as shown below?
CaCO3 + 2HCl
CaCl2 + H2O +
CO2
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
…………………………………………………………………………………………..
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
16.
When pure zinc reacted with dilute sulfuric acid, 2.4 dm3 of hydrogen gas
were collected at r.t.p. Calculate the mass of zinc?
Zn +
H2SO4
ZnSO4 + H2
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
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IGCSE
17.
Grade (10)
3
What volume of 0.4 mol/dm Hydrochloric acid is needed to react completely
with 0.24 g of magnesium?
Mg
+
2 HCl
MgCl2 + H2
……………………………………………………………………………………………
…………………………………………………………………………………………..
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
18.
Calculate the mass of silver formed when 5.52 g of silver carbonate are
heated, and find the volume of gases produced?
2 Ag2CO3
4 Ag + 2 CO2 + O2
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
…………………………………………………………………………………………….
……………………………………………………………………………………………
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IGCSE
19.
Grade (10)
Find the volume of ammonia gas formed when 1.605 g of ammonium
chloride is heated?
NH4Cl(s)
NH3 (g) + HCl (g)
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
…………………………………………………………………………………………….
……………………………………………………………………………………………
…………………………………………………………………………………………….
20. In an experiment, 25.0 cm3 of aqueous sodium hydroxide, 0.4 mol/dm3 was
neutralized by 20.0 cm3 of aqueous oxalic acid, H2C2O
2NaOH + H2C2O4
Na2C2O4 + 2 H2O
Calculate the concentration of the oxalic acid in mol/dm.
(a)
Calculate the number of moles of NaOH in 25.0 cm3 of 0.4 mol/dm3 solution
……………………………………………………………………………………………
…………………………………………………………………………………………….
……………………………………………………………………………………………
(b)
Use your answer in (a) to find out the number of moles of H2C2O4 in 20 cm3 of
solution.
……………………………………………………………………………………………
………………….................................................................................................................
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IGCSE
Grade (10)
3
3
(c) Calculate the concentration of the 0.04 mol/dm NaOH (aq) in g/dm
……………………………………………………………………………………………
……………………………………………………………………………………………
(c)
Calculate the concentration, mol/dm3 of the aqueous oxalic acid
……………………………………………………………………………………………
………………….................................................................................................................
……………………………………………………………………………………………
…………………………………………………………………………………………….
21. How many moles of Cu are required for the production of 15 moles of water
for the following:
[Cu(H2O)5] 2+
Cu2+
+
5H2O
………………………………………………………………………………………….
…………………………………………………………………………………………..
…………………………………………………………………………………………..
………………………………………………………………………………………….
…………………………………………………………………………………………..
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IGCSE
Grade (10)
Percentage purity:
Calculated mass of reactant
% purity
=
---------------------------------------- X 100
Given mass of reactant
22.
When 10 g of impure zinc reacted with dilute H2SO4, 2.4 dm3 of hydrogen
gas ware collected at r.t.p. Calculate the percentage purity of zinc?
Zn
+
H2SO4
ZnSO4 + H2
……………………………………………………………………………………………
…………………………………………………………………………………………….
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
…………………………………………………………………………………………….
……………………………………………………………………………………………
23.
2 g of an alloy (Cu and Al) reacted with HCl, 2.4 dm3 of H2 evolved.
Calculate the % purity of Al in the alloy?
2Al + 6 HCl
2 AlCl3 + 3 H2
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
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IGCSE
Grade (10)
Percentage Yield :
Actual mass
% Yield
=
--------------------------- X 100
Calculated mass
24.
Heating 12.4 g of copper (II) carbonate produced only 7 g of copper (II)
oxide. What is the % yield of copper oxide?
CuCO3
CuO + CO2
……………………………………………………………………………………………
…………………………………………………………………………………………….
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
Percentage composition of element in a compound :
Number of atoms of element X its Ar
Percentage composition = ---------------------------------------------
X 100
Mr of the compound
25. Find the % composition of nitrogen In the following :
(a)
Ammonium nitrate
.……………………………………………………………………………………………
…………………………………………………………………………………………….
(b)
Ammonium sulfate
……………………………………………………………………………………………
……………………………………………………………………………………………
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IGCSE
Grade (10)
26. Calculate the percentage composition of hydrogen and oxygen in H2O.
……………………………………………………………………………………
……………………………………………………………………………………
……………………………………………………………………………………
27.Calculate the percentage composition of H2O in CuSO4. 5H2O & the percentage
composition of oxygen.
……………………………………………………………………………………………..
……………………………………………………………………………………………..
…………………………………………………………………………………………….
…………………………………………………………………………………………….
Limiting and excess:
28.
3 g of magnesium was added to 12.0 g of ethanoic acid
Mg
+
2 CH3COOH
(CH3COO)2 Mg +H2
The mass of one mole of Mg is 24 g.
The mass of one mole of acid is 60 g.
Which one, magnesium or ethanoic acid, is an excess? Show your
work.
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
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IGCSE
(a)
Grade (10)
How many moles of hydrogen were formed?
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
(b)
Calculate the volume of hydrogen formed measured at r.t.p.?
……………………………………………………………………………………………
……………………………………………………………………………………………
29. Calculate the mass of FeS produced from the reaction between 28g of iron
with 20g of sulfur
Fe + S
FeS
……………………………………………………………………………………………
…………………………………………………………………………………………….
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
30. Calculate the mass of CO2 produced from the reaction between 2 dm3 of
C2H4 with 0.5 dm3 of O2 at r.t.p.
……………………………………………………………………………………………
………………………………………………………………………………………….
……………………………………………………………………………………………
………………………………………………………………………………………….
…………………………………………………………………………………………
………………………………………………………………………………………….
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IGCSE
Grade (10)
Empirical & molecular formulae:
C6H6
=
CH is the empirical formula
(CH) n
n= 6
C6H6 is the molecular formula
Empirical formula shows the simplest whole number ratio of atoms present in the
compound.
Molecular formula shows the number and type of atoms forming the molecule of the
compound
Mr of the molecular formula
n = …………………………………………………
Mr of the empirical formula
N.B. Hydrocarbon: made of C, H only
Carbohydrate: made of C, H & O
31.
A hydrocarbon contains 80 % carbon. Find the empirical formula and the
molecular formula, if the Mr of the compound is 30.
……………………………………………………………………………………………
………………………………………………………………………………………
……………………………………………………………………………………………
…………………………………………………………………………………………….
……………………………………………………………………………………………
……………………………………………………………………………………………
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IGCSE
Grade (10)
32. A carbohydrate has 40% of its mass carbon, 6.66% hydrogen. Find the
empirical and molecular formulae, given that its Mr is 180.
……………………………………………………………………………………………
……………………………………………………………………………………………
…………………………………………………………………………………………….
……………………………………………………………………………………………
……………………………………………………………………………………………
33. Calculate the empirical formula of the formed compound from the reaction of
0.24g Mg with 0.16g oxygen.
…………………………………………………………………………………………
…………………………………………………………………………………………
…………………………………………………………………………………………
………………………………………………………………………………………….
34. Compound contains 0.12g carbon & 0.02 g hydrogen. Calculate the empirical
formula of this compound; calculate the molecular formula of the compound if
its formula weight is 56g.
……………………………………………………………………………………
……………………………………………………………………………………
……………………………………………………………………………………
……………………………………………………………………………………
……………………………………………………………………………………
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IGCSE
Grade (10)
35.Calculate the empirical formula of an organic compound containing 92.3%
carbon & 7.7% hydrogen by mass. If the molecular weight of the organic
compound is 78, what is its molecular formula?
……………………………………………………………………………………
……………………………………………………………………………………
……………………………………………………………………………………
……………………………………………………………………………………
…………………………………………………………………………………………..
36. 4.2g of cerium reacted with oxygen to form 5.16g of an oxide of cerium.
Complete the following to determine the formula of this oxide.
Number of moles of cerium atoms used ………………………………………….
Mass of oxygen that reacted………………………………………………………..g
Number of moles of oxygen atoms in oxide………………………………………
Ratio by moles of cerium atoms to oxygen atoms………………………………….
Formula of the oxide of cerium is…………………………………………………..
37. The Mr of oxalic acid is 90 and its composition by mass is:
Carbon = 26.7%
hydrogen = 2.2%
oxygen = 71.1%
(I) Calculate the empirical formula of oxalic acid.
……………………………………………………………………………………
……………………………………………………………………………………
……………………………………………………………………………………
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IGCSE
Grade (10)
(II) What is the molecular formula of the acid?
………………………………………………………………………………………….
………………………………………………………………………………………….
…………………………………………………………………………………………
38. An excess of hydrochloric acid was added to 1.23g of barium carbonate. The
volume of carbon dioxide collected at r.t.p. was 0.120dm3. The impurities did
not react with the acid. Calculate the percentage purity of the barium
carbonate.
BaCO3 + 2HCl → BaCl2 + CO2 + H2O
Molar gas volume at r.t.p. is 24 dm3.
(i) The numbers of moles of CO2 collected ………………………….. …………mole
(ii) The numbers of moles of BaCO3 reacted ……………………….……………mole
(iii) Mass of mole of BaCO3…….………………………………………..……………g
(iv) Mass of barium carbonate …………………………………………...……………g
(v) Percentage purity of barium carbonate ……………………………..……………
40. 2g (an excess) of iron is added to 50 cm3 of 0.5 M sulfuric acid. When the
reaction is over, the reaction mixture is filtered. The mass of the unreached iron
is found to be 0.6g. (Fe = 56)
a. What mass of iron took part in the reaction?
……………………………………………………………………………………………
b. How many moles of iron atoms took part?
……………………………………………………………………………………………
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IGCSE
Grade (10)
c. How many moles of sulfur acid reacted?
……………………………………………………………………………………………
d. Write the equation for the reaction, and deduce the charge on the iron ion that formed.
……………………………………………………………………………………………
e. What volume of hydrogen (calculated at rtp) bubbled off during the reaction?
……………………………………………………………………………………………
41. When calcium carbonate is heated strongly, this chemical change occurs:
CaCO3 (s)
→ CaO (s) + CO2 (g)
a. Write a word equation for the change.
…………………………………………………………………………………………
b. How many moles of CaCO3 are there in 50g of calcium carbonate?
…………………………………………………………………………………………
c. (i) What mass of calcium oxide is obtained from the thermal decomposition of 50g
of calcium carbonate, assuming a 40% yield?
……………………………………………………………………………………………
(ii) What mass of carbon dioxide will be given off at the same time?
……………………………………………………………………………………………
(iii) What volume will this gas occupy at r.t.p. ?
…………………………………………………………………………………………….
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IGCSE
Grade (10)
Topic 5
The Periodic Table
­ The periodic table is a way to classify the elements
­ The elements are arranged according to the increase of their proton number where
elements show periodicity.
Group: is a set of elements are arranged vertically, having same valence electrons,
valency and consequently chemical properties but different number of energy
levels.
­ Group number indicates valence electrons.
­ Period: is a set of elements are arranged horizontally having same number of energy
levels but different valence electrons, valency and chemical properties [7 periods].
­ Period number indicates number of energy levels (electron shells).
­ The heavy zig-zag line separates metals from nonmetals
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IGCSE
Grade (10)
­ Metals are found at the left of the zig­zag line.
­ Non-metals are found on the right of the zig­zag line.
­ A group of elements are called “metalloids” [semiconductors], separate the metals
from the non­metals.
­ Hydrogen sites alone, as it has one valence electron but chemically behaves as
nonmetal. It has unique properties.
­ Inert gases are found at the far right of the table (group 0 or group VIII).
­ Starting from the fourth period, at the middle, transition metals are found.
­ Metallic properties decrease across the period, and increase down the group.
­ Non­metallic properties increase across the period and decrease down the group.
­ Alkaline properties decrease across the period, and increase down the group.
­ Acidic properties increase across the period, and decrease down the group.
­ Differences between metals and non­metals.
­ Artificial elements (created in lab) mostly are in the lowest block in the bottom row.
They are radioactive and their atoms break down very quickly (That is why they are
not found in nature)
­ Now if you know where an element is, in the periodic table, you can use the pattern
and trends to predict how it will behave.
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Grade (10)
Differences between
Metals and Non-metals
• Physical Differences:
Metals
Usually have high melting and boiling
points. Solids at room temperature.
Exceptions:
Group I metals have low melting
points and mercury is a liquid.
Good conductors of both heat and
electricity.
Hard
Exceptions:
Group I metals are soft.
High densities
Exceptions:
Group I metals have low densities.
Malleable (can have their shape
changed by hammering)
Ductile (can be pulled into wires)
Sonorous (think of a bell)
Can be polished to a luster (shiny)
Non-metals
Melting points and boiling points are low.
Exceptions:
Those non­metals with macromolecular
structures – carbon, silicon, boron, germanium.
Poor conductors
Exception: graphite
Soft
Exceptions:
Those non­metals with macromolecular
structures – carbon, silicon, boron, germanium.
Low densities.
Brittle (do not change shape but break)
Exception: diamond.
Not sonorous
Have a dull surface
Exceptions: graphite, iodine, diamond.
• Chemical Differences:
Chemical
property
Metals
Non-metals
Electron distribution
and bonding
1, 2 or 3 valency electrons.
These are lost to form cations
(positive ions).
Form ionic compounds with
non­metals
Many react with dilute acids to
give a salt and hydrogen.
4, 5, 6 or 7 valency electrons.
Either gain electrons to form anions
(negative ions) or share electron pairs.
Form covalent compounds with other
non­metals
Do not react with acids to give a salt
and hydrogen.
Type of compound
Reaction with acids
67
IGCSE
Grade (10)
Group properties
“Group I [Alkali Metals]”
• React with water forming alkalis, (pH 12
14).
Physical properties:
Down the Group
­
­
­
­
­
­
The alkali metals are not typical metals.
Soft and the softness increase [Li is the hardest one].
Melting point and boiling point decrease.
Density increases [Na, K are out of step, k has odd denisty].
Low density [first 3 float on water].
Good conductors of heat and electricity.
Shiny surface.
Chemical properties:
­ Monovalent.
­ Reactivity increases down the group, as loss of electrons is much easier.
­ Since they are highly reactive, they are stored under oil.
­ They form white solid compounds dissolve in water forming colourless solutions.
68
IGCSE
Grade (10)
1. Reactions with water
Form metal hydroxides (alkalis) and hydrogen gas evolved.
Metal + Water
Metal hydroxide + Hydrogen
2Li
+
2H2O
2LiOH
+
H2
N.B Exothermic reaction and becomes explosive as we go down the group.
♣ Li
Readily reacts, floats, melts, bubbles
♣ Na
Vigorous, floats, shots across the water surface, fizzing may catch fire
[golden yellow]
♣ K
Violent, floats, effervescence, hissing sound, catches fire [lilac flame]
♣ Rb
Explosive
Precautions taken when group (I) metal is put in water:
• Eye goggles.
• Behind a screen [fume cupboard].
• Small pieces of metal in large amount of water.
* All metals are reducing agents as they lose electrons during oxidation reactions
forming positive ions * their oxides are basic.
2. Reaction with chlorine:
Heating the three metals, and plunging them in gas jars of chlorine, they burn brightly
forming chlorides.
2Na(s) + Cl2 (g)
2NaCl(s)
3. Reaction with oxygen
Heating the three metals, and plunging them in gas jars of oxygen, they burn fiercely
forming oxides.
Uses:
♣ Li
♣ Na
♣K
Batteries
Street lamps
Fertilizers
69
IGCSE
Grade (10)
“Group VII [Halogens]”
Physical properties:
Down the group
1. Non­metals.
2. Diatomic.
3. Poisonous.
4. Melting point, boiling point, and density increase.
5. Bad conductors of heat and electricity.
6. The colour is getting darker
Halogen
Fluorine
At room temperature the element is...
F2
Boiling point/° C
Pale yellow gas
­188
Chlorine Cl2
Yellowish green gas
­35
Bromine Br2
Reddish brown liquid
59
Iodine
Black solid
184
I2
Chemical properties:
1. Oxidizing agents as they gain electrons during reduction reactions forming
negative ions
2. Monovalent.
3. React with metals forming salts
2Na + Cl2
2NaCl
2Fe + 3Br2
2FeBr3
2Al + 3I2
2AlI3
4. Chemical reactivity decreases down the group.
Cl2 (g) + 2KBr (aq)
2KCl (aq) + Br2 (aq)
Colourless
Cl2 (g) +
2KI (aq)
Orange
2KCl (aq)
Colourless
Br2 (l) +
2NaI (aq)
+ I2(aq)
dark brown
2NaBr (aq) + I2(aq)
Colourless
dark brown
N.B. A halogen will displace a less reactive one from a solution of its halide.
All group I metal halides are colourless.
This type of reaction is displacement and redox reaction.
70
IGCSE
Grade (10)
Transition Metals
Metals are found at the middle of the periodic table.
Physical properties:
­ High density.
­ High melting and boiling point.
­ Hard as their atoms are compacted.
Chemical properties:
­ Less reactive, do not corrode readily in the atmosphere. But iron is an exceptional
case­it rusts easily.
­ They show no clear trend in reactivity.
­ Do not react with cold water.
­ Most of them react with steam forming their oxides.
FeO(s) + H2(g)
Fe(s) + H2O(g)
Iron
­ Have more than one oxidation state. The Roman numeral tells its oxidation state.
­ Form coloured compounds except zinc.
• Used as catalysts:
Ex. Fe
in making ammonia by Haber process.
V2O5
in contact process which is a step in manufacture of H2SO4
71
IGCSE
Grade (10)
Redox
Oxidation and reduction reactions always take place at the same time.
Oxidation
­ Gain of O2
­ Loss of H2
­ Loss of electrons (OIL)
Reduction
­ Loss of O2
­ Gain of H2
­ Gain of electrons (RIG)
Reduction
CuO
+
H2
Oxidizing
agent
Reducing
agent
Cu + H2O
Oxidation
Oxidation
3H2 + N2
2NH3
Reduction
Reduction
Fe2O3 + 3CO
2Fe + 3CO2
Oxidation
Question:
Detect the oxidizing and reducing agents?
2H2 + O2
2H2O
If a substance loses electrons during chemical reaction, it has been oxidized.
If it gains electrons, it has been reduced. The reaction is a redox reaction
N.B. respiration, rusting, and burning are redox reactions
72
IGCSE
Grade (10)
Writing half equations to show the electron transfer:
2Mg + O2
2MgO
Half equation
2 Mg
2Mg2+ + 4e­
[Oxidation]
Half equation
O2 + 4e­
2O2­
[Reduction]
Number of lost electrons must equal number of gained electrons
From half equation to the ionic equation:
Ionic equation
2Mg
2Mg2+ + 4e­
O2 + 4e­
2Mg + O2
2O2­
2Mg2+ + 2O2­
2Na + Cl2
2NaCl
Example
Half equations:
Ionic equation
2Na
2Na+ + 2e­
Cl2 + 2e­
2Cl­
_______________________________
2Na+ Cl2
2Na+ + 2Cl­
Question: Write half equations and ionic equation for the following displacement
reaction?
Cl2 + 2KBr
2KCl + Br2
• Oxidation state: is the number of (+ve) or (–ve) charges of the ion in a compound.
­ The oxidation state is always given in as a roman numeral.
­ Number
1
2
3
4
5
6
7
­ Roman numeral
I
II
III
IV
V
VI VII
­ If oxidation state changes during a reaction, it is a redox reaction.
Oxidation is the increase in oxidation number
­IV
­III
–II
­I
0
+I
+ II
+III
+IV
Reduction is the decrease in oxidation number
73
IGCSE
Grade (10)
Oxidation number rules:
1. The oxidation number of any uncombined element is zero, ex. Zn, O2, S8.
2. In compounds many atoms or ions have fixed oxidation numbers
­ Group I elements are always +1
­ Group II elements are always +2
­ Hydrogen is +1 ( except in metal hydrides such as NaH when it is ­1)
3. The oxidation number of an element in a monoatomic ion is always the same as
the charge, ex. Cl1­ is ­1, and Al3+ is +3.
4. The sum in oxidation numbers in a compound is zero.
5. The sum of the oxidation numbers of ions in a compound ions is equal to the
charge on the compound ion.
State the oxidation number of the bold atom in these compounds or ions:
1. SO2
2. ICl3
3. SO42­
4. ClO24Fe + 3O2
2 Fe2O3
O.S = zero
O.S = +3
Oxidation
4Fe3+
6 O2-
4 Fe
6 O + 12e-
+ 12 e-
[Oxidation]
[Reduction]
Oxidation
Mg
0
CuO + Mg
Cu2+
Mg2+
Cu + MgO
Cu0
Reduction
N.B. In electrolysis:
* [At anode]
e– loss
(oxidation)
* [At cathode]
e– gain
(reduction)
74
IGCSE
Grade (10)
Colour changes in redox reactions [change in oxidation state]:
1. Acidified potassium dichromate VI is an oxidizing agent (K2 Cr2O7)
The oxidation state changes from +6 to +3
Reduction
+6
2Cr3+
Cr
Orange
Green
2. Acidified potassium manganate VII (KMnO4) is an oxidizing agent with a purple
colour. The oxidation state changes from +7 to +2 which is more stable, by
gaining electrons (reduction).
Reduction
7+
Mn
Mn2+
Purple
colourless
If a reducing agent is present, the purple colour will fade.
3. Iodine is an oxidizing agent when reduced to iodide.
I2 + 2 e ­
Dark brown
2 I­
Colourless
4. Potassium iodide is used to test for the presence of an oxidizing
agent as hydrogen peroxide
H2O2(aq) + 2KI(aq) + H2SO4(aq)
I2(aq) + K2SO4(aq) + H2O(aq)
Oxidation
-
2I (aq)
I2 (aq)
Colourless
Dark brown
5. From the famous reducing agents are H2, CO, SO2, C and KI. All can be used to
test for oxidizing agent.
Remember:
• All nonmetals are oxidizing agents [strongest two are F and Cl]
• All metal ions and H+ are oxidizing agents
• All metals are reducing agents
[strongest one is cesium]
• All nonmetal ions are reducing agents.
75
IGCSE
Grade (10)
Topic 6
Electrochemistry
“Conductors”
• Substances that allow electricity to pass through easily.
• Ex. - All metals and graphite, as they have mobile (freely moving) valence electrons.
-Molten or solution form of ionic compounds as they have freely moving ions.
+
Ammeter
–
Bulb
A
Metal
♣ Compare the electrical conductivity of two metals?
- The electric circuit is set up, where the battery acts like electron pomp
- Electrons move from the negative terminal to the positive one through the circuit.
- Ammeter reading is recorded.
- For fair test use the same length and cross thickness of wire for each metal.
Uses of Conductors:
Copper:
Is used in electrical wiring as:
1. It is good conductor of electricity.
2. Ductile (easily drawn into wires).
3. Easily purified.
Aluminium:
Is used in making high-voltage power lines as:
1. It is good conductor of electricity.
2. Resists corrosion.
3. Has low density.
N.B. A steel core can be used to give the high voltage cables additional strength to
stop them sagging and breaking. Aluminium is better conductor than steel.
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Grade (10)
“Insulators”
• Substances that resist flow of an electric current.
• Ex. Plastics, glass and ceramics as they do not have mobile electrons also all nonmetals except graphite.
Uses of Insulators:
Plastics:
Ex. PVC (poly vinyl chloride).
Used to cover electric wires for safety as they are :
1. Bad conductors of electricity, to avoid electric shocks
2. Flexible, easily molded.
3. Non-biodegradable (do not decayed by the effect of bacteria).
4. Cheap, recycled.
Ceramics:
Made by heating clay.
Used in making fuses, base of electric iron and electric heaters as:
1. They do not conduct electricity.
2. Have high melting point, used in high temperature.
[Used in high-voltage electricity towers (pylons) to keep these wires
away from touching each other and electricity from running down the
pylon-dangerous]
3. Not affected by water or oxygen of air.
4. Can be molded to complex shapes.
77
IGCSE
Grade (10)
“Electrolysis”
• is the decomposition of an ionic compound, when molten or in solution by passage of
an electric current
• is a decomposition reaction, as a chemical compound breaks down into simpler
substances
• Ionic compounds conduct electricity when molten or dissolved in water as they have
free moving ions.
• Electrical energy changed to chemical energy [ endothermic]
DC-Power
Anode (+)
Cathode (–)
Molten
PbBr2
Electrodes
(Graphite)
Electrolytic Cell
Electrolytic cell consists of :
1. Two electrodes (cathode and anode)
2. External curcit with DC power source
3. Electrolyte molten or aqueous substance that undergoes electrolusis
• Electrolyte: is the compound that conducts electricity when molten or dissolved in
water and breaks down during electrolysis.
• Electrons move from the negative terminal of the battery to cathode, then from anode
to positive terminal of the battery.
• The electrodes are made of graphite or platinum (inert electrodes).
• Electrodes: are rods that carry the electric current to and from the electrolyte.
Electrolysis of molten lead bromide
N.B. This experiment is carried out in a fume cupboard because Pb & Br2 have toxic
vapours.
78
IGCSE
Grade (10)
[PbBr2]
Pb2+
Attracted to cathode
(–)
Pb2+ + 2e–
Pb
(Reduction)
Half equations
Observation:
- At cathode
- At anode
2Br–
Attracted to anode
(+)
2Br– – 2e–
Br2
(Oxidation)
grey colour of molten Pb
reddish brown gas of Br2
Predicting the Products of Electrolysis
Molten Ionic Compounds:
Compound
electrolyzed
Product at cathode
(–)
Product at anode
(+)
Aluminium oxide
Aluminium
Oxygen
Copper II bromide
Copper
Bromine
Sodium chloride
Sodium
Chlorine
Zinc II Iodide
Zinc
Iodine
79
IGCSE
Grade (10)
The rules for the electrolysis of a solution:
-
At the cathode (-), either a metal or hydrogen forms.
The going down the reactivity series, the more likely the ion will be discharged
[changed into atom or molecule at the electrode].
If the metal is more reactive than hydrogen, its ions stay in the solution and hydrogen
is discharged.
If the metal is less reactive than hydrogen, the metal forms.
The most reactive metals form the most stable ions; these ions will be difficult
to convert back to metals.
N.B. The least reactive element will discharge first.
At the anode (+), a non-metal forms.
- If the electrolyte is a concentrated solution of halide (Cl-, Br- or I-), then chlorine,
bromine or iodine is formed.
- If the halide solution is dilute, or there is no halide, oxygen forms.
Ease of discharge
K+
Na+
Mg2+
Al3+
H+
Cu2+
Ag+
SO42–
NO31–
OH–
Cl–
Br–
I–
These ions
never
discharged
80
IGCSE
Grade (10)
“Solutions of Ionic Compounds”
Concentrated sodium chloride solution (brine):
Brine is obtained by evaporating sea water, till we get concentrated solution.
81
IGCSE
Grade (10)
NaCl
Na+
-
H2 O
Cl–
H+
OH–
Half equations:
At cathode
2H+ + 2e–
H2
(colourless bubbles of gas)
At anode
2Cl– – 2e–
Cl2 (yellowish-green bubbles of gas)
[Na+, OH–] ions are left in the solution; some of the solution is evaporated to get a
more concentrated solution, or evaporated till dryness giving solid sodium hydroxide.
Concentration of NaOH is high at cathode
Remember half equation shows the electron transfer at an electrode.
The overall reaction is
2NaOH (aq) + Cl2 (g) + H2(g)
2NaCl (aq) + 2H2O (l)
What the products are used for ?
Chlorine is a poisonous – yellow green gas:
Used for making…..
1. The plastic PVC
2. Bleaching agent
3. Water treatment as antiseptic to kill bacteria
4. Weed killers and pesticides
5. Hydrochloric acid
Hydrogen is a colourless flammable gas:
Used for making…..
1. Rocket fuel and fuel for cars under experimental stages
2. Margarine
3. Hydrogen peroxide(H2O2)
4. Ammonia by Habber process
82
IGCSE
Grade (10)
Sodium hydroxide solution, alkaline and corrosive:
1. Soap
2. Detergents
3. Treatment of natural textile
4. Paper
5. Extraction of Aluminium
Diluted sodium chloride solution:
Try these solutions:
1) CuSO4 solution:
[blue]
CuSO4
Cu2+
SO42–
H2 O
H+
OH–
- At cathode
2Cu2+ + 4e–
2Cu
- At anode
4OH– – 4e–
O2 + 2H2O
- H2SO4 left (colourless), the blue colour disappeared.
83
IGCSE
Grade (10)
Electrolysis of water:
- Drops of H2SO4 must be added to conduct electricity as pure
water is a covalent compound [bad conductor of electricity].
H2 O
H+
OH–
- At cathode
4H+ + 4e–
2H2
- At anode
4OH– – 4e–
O2 + 2H2O
- H2SO4 is found in small quantity so it is left over, as water is
used up.
N.B.
Volume of H2 is doubled O2 as (2H & 1O)
[H2 : O2]
2:1
♣ Predict the products of electrolysis of:
1. Aqueous potassium iodide
2. Aqueous copper II nitrate
• Concentrated HCl:
HCl
H+
H2
(at cathode)
N.B.
Oxidation
Reduction
H2 O
Cl–
H+
OH–
Cl2
(at anode)
loss of electrons (at anode)
gain of electrons (at cathode)
Uses of Electrolysis:
1. Extraction of metals from their ores ex. Aluminium.
2. Purifying copper.
3. Electroplating.
84
IGCSE
Grade (10)
(1) Extraction of Aluminium
•
•
•
•
•
•
Aluminium is the most abundant metal in Earth’s crust
Aluminium ore is called “bauxite”.
The ore “bauxite” is purified from its impurities to form Alumina Al2O3.
Aluminium is extracted from aluminium oxide (Alumina).
Alumina melting point is ≅ 2000 °C which is very high, so cryolite is added.
The electrolyte is molten aluminium oxide and cryolite [Sodium aluminium fluoride
Na3AlF6].
•
Cryolite is added to:
1. Lower melting point of alumina to 900 °C
3. Helps conduction of electricity
Half equations:
• At cathode
4Al3+ + 12e–
• At anode
6O2– – 12e–
• The overall reaction is
2Al2O3 (l)
4 Al(l)
3O2(g)
2. Saves energy
Reduction
Oxidation
4Al (l) + 3O2(g)
85
IGCSE
Grade (10)
• Oxygen reacts with graphite rode (+) forming CO2 gas that escapes, so the
anode must be changed from time to time.
• Aluminium is used to make drinks cans, food cartons, cooking foil and aircraft.
(2) Electroplating
• Is used to plate one metal with a different one.
• The key must be cleaned by sand paper then cotton wool before electroplating so that
the copper sticks to the key.
• The key must be rotated during the electroplating to be covered from all direction.
Electroplating is used to:
1. Get beautiful (good shiny) appearance
Half equations:
• At anode
• At cathode
Cu – 2 e–
Cu2+ + 2 e–
2. Protects from corrosion.
Cu2+
Cu (key)
N.B.
Cu, Ni, Cr, Ag & Tin are the most metals used in plating.
The role of electrolyte is to:
* keep the concentration of Cu2+ constant
* conduct electricity
86
IGCSE
Grade (10)
(3) Refining Copper
• Copper is purified by electrolysis.
• At anode
• At cathode
Cu – 2e–
Cu2+ + 2e–
Cu2+ (decreases in size)
Cu (increases in size)
N.B.
The impurities reduce the electrical conductivity and increase electric resistance.
N.B. Electroplating and refining of copper take place by electrolysis with active
electrodes.
87
IGCSE
Grade (10)
Topic 7
Properties of Acids and Bases
“Acids”
• Acid: is a proton donor, reacts with base to form salt and water.
HCl(aq.)
H+(aq) + Cl–(aq)
• Acids are pure compounds in water and can be used as diluted or concentrated.
Physical properties of acids:
1. Sour taste.
2. Has corrosive effect.
3. pH is less than 7.
4. Turns litmus paper red, methyl orange red and thymolphthalein colourless.
• Mineral acids: are strong acids, completely ionized in water i.e. good proton donor
[good conductors of electricity because there more ions present].
• Their solutions have high concentration of H+ having low pH value [pH 1 3].
Ex. HCl , H2SO4 , HNO3
H2SO4 (aq)
2H+(aq) + SO42–(aq)
Good Conductor of electricity
• Organic acids: are weak acids, partially ionized in water i.e. weak proton donor [weak
conductors of electricity].
Ex. Lemon juice contains citric acid, ant stings contain methanoic acid, fizzy drinks
contains carbonic acid and vinegar which is Ethanoic acid (CH3COOH)
• Their solutions have low concentration of H+
[pH 4 6]
CH3COOH(aq)
CH3COO–(aq) + H+(aq)
acetic acid
Weak Conductor of electricity
• The higher the concentration of hydrogen ions, the lower the pH
88
IGCSE
Grade (10)
Chemical properties of acids:
1) Metal (active metal) + Acid
Salt + Hydrogen [Displacement/redox/ Exo.]
Mg(s) + 2HCl(aq.)
MgCl2(aq.) + H2(g)
Ionic Equation
Mg(s) + 2H+(aq)+ 2Cl–(aq)
Mg(s) + 2H+(aq)
Mg2+(aq)+ 2Cl–(aq) + H2(g)
Mg2+(aq) + H2(aq)
Redox reactions because electrons are transferred from Mg to H+ .
2) Metal oxide + Acid
Salt + Water
CuO(s) + H2SO4(aq)
Ionic Equation
CuO(s) + 2H+(aq)+ SO42–(aq)
CuO(s) + 2H+(aq)
[Neutralization/ Exo.]
CuSO4(aq.)
+ H2O(l)
Cu2+(aq) + SO42–(aq) + H2O(l)
Cu2+(aq) + H2O(l)
Neutralization reactions are not redox reactions as no change in oxidation state.
3) Metal hydroxide + Acid
Salt + Water
[Neutralization / Exo.]
NaOH (aq) + HCl(aq.)
NaCl(aq.) + H2O(l)
OH- (aq) + H+ (aq.)
H2O(l)
N.B. The base ammonia neutralizes acids to form ammonium salts
NH3 (g) + HNO3(aq)
NH4NO3 (aq)
NH3 (g) + HCl(g)
NH4Cl(s)
4) Carbonate + Acid
Salt + H2O + CO2
CaCO3(s) + H2SO4(aq.)
[Neutralization / Exo.]
CaSO4(s) + H2O(l) + CO2(g)
N.B
Acid
+ Metal
Salt + H2
+ Metal oxide or hydroxide
Salt + H2O
+ Metal carbonates
Salt + H2O + CO2
89
IGCSE
Grade (10)
“Bases”
Base: is a proton, acceptor reacts with acid to form salt and water.
A base can be metal oxide or hydroxide.
An alkali is a soluble base (aq.).
Some substances act as bases.
The pure alkalis are solids except ammonia which is a gas. They are used in lab as
aqueous solutions.
• Examples: aq. ammonia (NH4OH), metal carbonate (CaCO3), metal hydrogen
carbonates (NaHCO3).
• All react with acid to form salt and water.
Physical properties of alkalis:
1. pH is more than 7.
2. Soapy feeling.
3. Corrosive (concentrated).
4. pH is more than 7.
5. Turns both litmus paper and thymolphthalein blue, and methyl orange yellow.
•
•
•
•
•
• Strong alkalis: are completely ionized in water releasing high concentration of
hydroxide ions i.e. good proton acceptor [pH ≈ 12 14]
NaOH(aq)
Na+(aq) + OH–(aq)
Good Conductor of electricity
• Weak alkalis: are partially ionized in water releasing low concentration of hydroxide
ions i.e. weak proton acceptor [pH ≈ 8 10]
NH4OH (aq)
NH4+(aq) + OH–(aq)
Weak Conductor of electricity
N.B. 1. All hydroxides are water insoluble except group I hydroxides and calcium &
barium hydroxides.
2. The higher the concentration of hydroxide ions, the higher the PH
90
IGCSE
Grade (10)
Chemical properties of bases:
1) Metal oxide + Acid
CuO(s) + H2SO4 (aq.)
Salt + H2O
CuSO4 (aq.) + H2O(l)
Metal hydroxide + Acid
NaOH (aq.)
Salt
+ H2 O
+ HCl (aq.)
NaCl(aq.) + H2O(l)
Na+(aq) + OH–(aq)+ H+(aq) + Cl–(aq)
Na+(aq) + Cl–(aq)+ H2O(l)
OH–(aq) + H+ (aq.)
H2O (l)
2) Ammonium salt + Alkali
Salt + Ammonia + Water
NH4Cl(aq) + NaOH(aq)
NaCl(aq.) + NH3(g) + H2O(l)
NH4+(aq) + OH–(aq)
NH3(g) + H2O(l)
(Pungent smell)
N.B. Bases such as sodium, potassium and calcium hydroxide react with ammonium salts,
driving out ammonia gas with pungent smell.
Ca(OH)2(s) + 2NH4Cl(s)
CaCl2(s) + 2H2O(l) + 2NH3(g)
Indicators:
Indicators are used to tell if a solution is acidic, alkaline, or neutral
The indicator paper must be wet when used with gases, to allow the substance to be
dissolved, ionized and release H+ or OH–
Indicator
Litmus paper or solution
Methyl orange
phenolphthalein
Thymolphthalein
Colour in acid
red
red
colourless
colourless
Colour in Neutral
Purple
Orange
colourless
colourless
Colour in alkali
blue
yellow
pink
blue
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Thymolphthalein
pH Scale”
• Measures the degree of acidity or alkalinity in a solution.
** Scale from 0 to 14 **
1
• Measures the concentration of H+.
pH α
+
[H ]
N.B. pH can be measured accurately by using pH meter or pH paper
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Grade (10)
PH mater is used to measure the
concentration of H+ directly by dipping the
electrode in the solution. It is accurate as it
measure to decimal.
Universal indicator:
Universal indicator is a mixture of dyes used to determine the degree of acidity or
alkalinity of a solution. It can be used as a solution or a paper strip.
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“Types of Oxides”
Oxides are compounds containing oxygen and another metal.
Elements
Metals
Non-metals
Oxygen
Oxygen
Metal Oxide (ionic) Solids
Basic oxide
Soluble
Ex. KOH ,
Ca(OH)2 ,
NaOH
(Alkalis)
Insoluble
Ex. CuO ,
FeO
Non-metal Oxide [covalent]
“Most are gases”
Amphoteric
Ex. Al2O3 ,
ZnO
Most are acidic
Ex. CO2 ,
NO2 , SO2 ,
SO3
Neutral
Ex. NO,
N2O,
CO, H2O.
• Magnesium ribbon burns with white flame leaving a white ash of magnesium oxide
• Hot iron reacts with oxygen. It glows bright orange and throws out a shower of
sparks. Black iron oxide is left.
• Copper is too unreactive to catch fire in oxygen. But when is heated in a stream of the
gas, its surface turns black copper oxide.
• Hot powdered carbon reacts with oxygen. It glows bright red and carbon dioxide is
formed, which is slightly soluble in water.
• Sulfur catches fire and burns with blue flame, and sulfur dioxide is formed.
• Phosphorous bursts into yellow flame in air or oxygen without heating (so it is stored
under water) forming a white solid Phosphorous (V) oxide.
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Basic oxides:
Metal oxides react with acids to form salt and water.
MgO(s) + 2HCl(aq)
MgCl2(aq) + H2O(l)
White
colourless
Amphoteric oxide is a metal oxide that shows both acidic and basic properties. They
react with both acids and bases forming salt and water.
Ex. Al2O3 & ZnO
Al2O3(s) + 6HCl(aq)
2AlCl3(aq) + 3H2O(l)
Al2O3(s) + 2NaOH(aq)
2NaAlO2(aq) + H2O(l)
Sodium Aluminate
2 ZnO(s) + 4HCl(aq)
2ZnCl2(aq) + 2H2O (l)
ZnO (s) + 2NaOH(aq)
Na2ZnO2(aq) +
H2O (l)
Sodium Zincate
Acidic oxides:
Most of non-metal oxides are acidic oxides as:
1. They dissolve in water forming acids.
SO3(g) + H2O(l)
H2SO4
2. They react with alkalis forming salt and water.
CO2(g) + 2NaOH(aq)
Na2CO3(aq)
+ H2O(l)
CaO(s) + SiO2(s)
CaSiO3(s)
Base
acidic oxide
Calcium silicate
Neutral oxides:
Some oxides of nonmetals are neutral; do not react with acids or bases.
H2O, CO, NO and N2O are neutral gases.
Dinitrogen oxide (N2O) is used as an anesthetic by dentists (laughing gas)
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Preparation of Salts
Salt is an ionic compound, can be prepared by reacting acids with metals, or insoluble
bases, or soluble bases (alkalis), or carbonates.
“Ionic Compounds”
Salts
+ve ion (cation)
- M+(metal ion)
- NH4+
–ve ion (anion)
- F– , Cl– , Br– , I–
- NO3–
- SO42–
- CO32–
• Solubility rules of salts
N.B.
1. All nitrates are soluble.
2. All sulfates are soluble except Calcium, Lead, and Barium.
3. All halides are soluble except Silver and Lead.
4. All carbonates are insoluble except sodium, potassium and ammonium carbonate.
To Prepare a Soluble Salt
Insoluble substance
Metal, Metal oxide,
hydroxide or carbonate
Acid
Excess method
Soluble
salt
Alkali (Titration)
Na+, K+, Ca+, NH4+
hydroxide
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To prepare magnesium sulfate crystals:
Metal + Acid
Mg(s) + H2SO4(aq.)
-
Grade (10)
Salt + Hydrogen
MgSO4(aq.) + H2(g)
Place a suitable amount of sulfuric acid in a beaker.
Add a piece of magnesium till fizzing stops and solid magnesium is seen.
Filter to remove the excess magnesium.
Heat the filtrate till point of crystallization, leave to cool, filter & dry between two
filter papers
violent reaction
N.B. 1. Na, K & calcium cannot be used as they are active metals
and the salt will not be pure as excess metal react with water forming metal hydroxide,
which is water soluble.
2. Water of crystallization is the water molecules present in hydrated crystals.
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To prepare copper II sulfate crystals:
CuO(s) + H2SO4(aq.)
Grade (10)
CuSO4(aq.) + H2O(l)
- Place a suitable amount of sulfuric acid in a beaker.
- Add copper oxide powder till the colour changes to blue and solid copper oxide is
seen.
- Filter to remove the excess copper oxide.
- Heat the filtrate till point of crystallization, leave to cool, filter & dry between two
filter papers.
- Blue crystals of hydrated copper sulfate CuSO4.5H2O
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Grade (10)
To prepare sodium chloride solution: [Titration]
Acid + Alkali
Salt + Water
HCl + NaOH
NaCl + H2O
- Using a pipette put 25 cm3of dilute sodium hydroxide in a conical flask, then add
few drops of phenolphthalein indicator, the solution turns pink.
- Fill in a 50 cm3 burette with dilute hydrochloric acid.
- Add the acid from the burette to the conical flask slowly a little at a time
- Shake the flask till the colour changes from pink to colourless, turn off the tap and
record the volume of acid.
- Repeat the experiment under same conditions [concentration, volume and
temperature] without using indicator; pour the solution in an evaporating dish.
- Heat the solution till point of crystallization, leave to cool, filter & dry between two
filter papers
Use of indicator:
To show that the reaction is completed as both reactants are soluble, colourless and no
bubbles of gas are formed.
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Finding concentrations by titration:
- Concentration of an acid can be calculated using a solution of an alkali of known
concentration (a standard solution) and titrate the acid against it.
N.B. Concentration is usually given in mol/dm3 or M.
1000 cm3 = 1 dm3
An example:
You are asked to find the concentration of a solution of sulfuric acid, using a 1 M
solution of sodium hydroxide as the standard solution.
- First, titrate the acid against standard solution.
- Calculate the volume of acid used
Starting volume of acid from the burette = 1.0 cm3
Final volume
= 28.8 cm3
Volume of acid used
= 27.8 cm3
So 27.8 cm3 of the acid neutralized 25 cm3 of the alkaline solution.
- Calculate the number of moles of sodium hydroxide used.
1000 cm3of 1M solution contains 1 mole so
25 cm3 contains 25 x 1 mole = 0.025 mole
1000
- From the equation, find the molar ratio of acid to alkali
H2SO4 (aq) + 2 NaOH(aq)
Na2So4(aq) + 2H2O(l)
1 mole
2 moles
?
0.025 moles
- Work out the number of moles of acid neutralized.
n = 0.025 x 1 = 0.0125 moles of acid were neutralized.
2
- Calculate the concentration of the acid
The volume of acid used is 27.8 cm3 = 0.0278 dm3
n = CxV
C = n/V
Concentration of acid = 0.0125 = 0.45 mol/dm3
0.0278
The concentration of hydrochloric acid is 0.45 M
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Preparation of insoluble salts by precipitation
• Precipitation is the formation of insoluble salt when two soluble salts react
together.
• A precipitate is an insoluble product of a reaction
• To precipitate an insoluble salt, we mix a solution that contains its positive ions
with one that contains its negative ions.
Ag NO3 ( aq) +
Ionic equation
NaCl (aq)
NaNO3 (aq)
Cl1-(aq)
AgCl(s)
Na2SO4 (aq)
2NaI (aq)
Ag1+ (aq) +
BaCl2 (aq
+
Pb (NO3)2 (aq) +
+
AgCl(s)
White ppt.
2 NaCl(aq)
+
BaSO4 (s)
White ppt.
2NaNo3 (aq)
+
PbI2 (S)
Yellow ppt.
Then filter to get the precipitate and dry between filter papers.
Some uses of precipitation:
1. Used to make coloured pigments for paints.
2. Used to remove harmful substances dissolved in water when cleaning up
waste water.
3. Used in making photographic films.
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Tests for gases
gas
ammonia (NH3)
carbon dioxide
(CO2)
chlorine (Cl2)
hydrogen (H2)
oxygen (O2)
sulfur dioxide
(SO2)
test and test result
turns damp red litmus paper blue
turns limewater milky
bleaches damp litmus paper
“pops” with a lighted splint
relights a glowing splint
turns aqueous potassium dichromate (VI) from orange to
green
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Flame tests:
• Dip a clean platinum or nichrome wire into conc. HCl and hold it in a hot
Bunsen flame.
• Dip the wire into the acid again, and then dip it into the salt, so that some
sticks to it.
• Hold it in the clear part of a blue Bunsen flame, and observe the colour.
Cation
Lithium (Li+)
Sodium (Na+)
Potassium (K+)
Calcium ( Ca2+)
Barium ( Ba2+)
Copper(Cu2+)
Flame colour
red
Orange - yellow
lilac
Orange - red
Light green
Blue - green
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QUALITATIVE ANALYSIS
Test for anions
Anion
Carbonate
(CO32–)
Chloride (Cl–)
[in solution]
Bromide (Br-)
[in a solution]
Iodide (I–)
[in solution]
Test
Add dilute acid
Acidify with dilute nitric acid,
then add aqueous silver nitrate
Acidify with dilute nitric acid,
then add aqueous silver nitrate
Acidify with dilute nitric acid,
then add aqueous lead (II)
nitrate
–
Nitrate (NO3 ) Add aqueous sodium hydroxide
[in solution]
then aluminium foil; warm
carefully
2–
Sulfate (SO4 ) Acidify with dilute nitric acid
[in solution]
then add aqueous barium
nitrate
Sulfite (SO32-)
- Add hydrochloric acid, then
In a solution.
heat
- Acidify with dilute nitric acid
then add aqueous barium
nitrate
Test result
Effervescence of carbon
dioxide
White precipitate (AgCl).
Creamy (off white)
precipitate (AgBr).
Yellow precipitate (PbI2).
Pungent smell of ammonia
(NH3)
White precipitate (BaSO4)
insoluble in excess.
- Colourless gas evolved
(SO2) changes purple aq.
acidified potassium
manganate (VII) paper
colourless
- White precipitate soluble in
excess acid.
N.B.
Sulfates are salts of sulfuric acid H2SO4
Sulfites are salts of sulfurous acid H2SO3
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Test for aqueous cations
Cation
Effect of aqueous sodium
hydroxide
Aluminium (Al3+) white ppt., soluble in excess
giving a colourless solution
2+
Zinc (Zn )
White ppt., soluble in excess
giving a colourless solution
Ammonium
Pungent smell of ammonia on
(NH4+)
warming
2+
Calcium (Ca ) white precipitate, insoluble in
excess
light blue precipitate, insoluble
Copper(II)
2+
in excess
(Cu )
Iron(II) (Fe2+)
Iron(III) (Fe3+)
Chromium
(Cr3+)
Green precipitate, insoluble in
excess
Red-brown precipitate,
insoluble in excess
Grey- green ppt. soluble in
excess giving dark green
solution
Effect of aqueous
ammonia
White precipitate., insoluble
in excess
White ppt., soluble in excess
giving a colourless solution
–
No precipitate or very slight
white ppt.
Light blue precipitate,
soluble in excess giving a
dark blue solution
Green precipitate, insoluble
in excess
Red-brown precipitate,
insoluble in excess
Grey- green ppt. insoluble in
excess
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Deduction that can be made from a substance appearance or smell:
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Topic 8
Metals
• Chemical activity series: is the arrangement of elements in a descending order
according to the decrease of their chemical reactivity.
• The order of reactivity, based on the reaction with water and dil. hydrochloric acid.
Reactivity series
Metals in order
of reactivity
Potassium
Sodium
Calcium
Magnesium
Zinc
Iron
Reaction with
water or steam
React violently with
cold water to form
the hydroxide and
hydrogen.
Reacts quickly with
cold water
Reaction with dilute
hydrochloric acid
Dangerous, explosive
violence
Very vigorous reaction
to form the chloride
and hydrogen.
Very vigorous reaction
to form the chloride
and hydrogen.
Very slowly with
cold water but
burns in steam to
form its oxide and
hydrogen.
React when heated
in steam to form the
solid oxide and
React to form the
metal chloride and
hydrogen.
Do not react with
hydrogen
cold water.
Reduction of oxide
with carbon
-
Metal oxides above
zinc cannot be
reduced with carbon.
All metal oxides
starting from zinc can
be reduced by heating
with carbon to form
the metal.
* Hydrogen
Copper
Does not react with
cold water or steam
Metals below
hydrogen do not react
with dilute acid.
Reduced to copper.
* The non-metal hydrogen is included in the first column of the table to show that metals above it react
with dilute acid to form a salt and hydrogen, whereas metals below it do not react with dilute acid.
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Heat
- A reactive metal has a strong tendency to lose electrons and form ion.
- The more reactive the metal, the more stable its ion and its compounds are. They do
not break down easily.
- The more reactive metal replaces less reactive one in its salt solution
- More reactive metals which above hydrogen in reactivity series, replace hydrogen of
acids
- Reaction with oxygen
K
Na Ca Mg Al
Burn very
brightly and
vigorously
C
Zn
Burn to form oxide
with decreasing vigor
Fe
Sn Pb
H Cu Hg Ag Au
React very
slowly to form
the oxide
Do not
react
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Making use of the reactivity series:
Extraction of metal from its ore:
- Highly reactive metals which are above C in chemical activity series are extracted from
their ores (highly stable compounds) by electrolysis. This is a powerful method, but it
costs a lot because it uses a lot of electricity.
- Less reactive metals which are below C in chemical activity series are extracted from
their ores by reduction using C. Many ores are oxides or compounds easily convert to
oxides which are reduced by carbon
Electrolysis
K
Na Ca Mg Al
Active metals
C Zn
Fe
Sn Pb
H Cu Hg Ag Au
Reduction by carbon
Native elements
- Unreactive elements are found in Earth’s crust native or uncombined.
Examples of metal extraction from their ores:
1. Extraction of iron from haematite
2. Extraction of aluminium from bauxite
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Extraction and Uses of Metals
“Extraction of Iron in Blast Furnace”
- Iron ore is haematite Iron III oxide.
- Blast furnace is fed (charged) through the top of the furnace by:
-
Haematite (Fe2O3), mixed with sand and other compounds
-
Limestone (CaCO3)
-
Coke (C), pure carbon, made from coal.
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Stage 1: The coke burns giving off heat
C(s)
+
O2(g)
CO2 (g)
(Exothermic step)
Redox
Stage 2: carbon monoxide is made
CO2(g)
+
C(s)
2CO(g)
Redox
Strong reducing agent
Fe2O3(s) + 3CO(g)
2Fe(s) + 3CO2(g)
Pig iron
Redox
CaO (s) + CO2(g)
Thermal decomposition
What is lime stone for?
CaCO3(s)
Lime
CaO(s)
Basic
oxide
+
SiO2(s)
acidic
oxide
Or CaCO3(s) + SiO2(s)
CaSiO3(s)
slag
less dense than Fe
Neutralization
CaSiO3(s) + CO2(g)
Pig iron is brittle and hard as it contains (4
5%) carbon from the coke.
Slag is used to reinforce the high ways and runways of airports.
The waste gases; hot carbon dioxide and nitrogen come out from the top of the furnace
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Uses of metals:
♣ Aluminium
Aircraft [low density]
Food container – Cooking pans (resists corrosion), nontoxic
High voltage cables [good conductor of electricity , low density].
Al resists corrosion due to the formation of a protective layer
Al2O3
♣ Zinc
Coins – Roofing alloys – Galvanizing of iron to resist corrosion.
♣ Copper
Wiring – Piping – Cooking pans
Strong and excellent conductor and easily drawn into wires and can
be bent easily
“Alloys”
- An alloy is a mixture of metal and one or more different element.
- Alloys are designed to have properties better suited for a particular use.
- Alloys are harder and more resistant to corrosion than the original metal
+
e–
+
e–
+
e–
+
e–
+
+
+
+
e–
e–
e–
e–
Presence of different sized atom
will prevent the layers
from slipping
The regular arrangement of a metal lattice structure is distorted in alloys
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♣ Brass
Grade (10)
Cu 70 %
Hard strong and shiny
Musical instruments,
Zn 30 %
♣ Bronze
door knobs and keys
Cu 90 %
Hard
Status, ornaments
hard
Car bodies, bridges, and
Tin 10 %
♣ Steel
Fe 99.7 %
C 0.3 %
♣ Stainless steel
construction of buildings
Fe 74 %
does not corrode
Kitchen sinks, Cutlery,
Cr 18 %
and chemical plants
Ni 8 %
♣ Solder
Tin 50 %
has lower melting point than its components
Lead 50 %
Making electrical connections
Aluminium alloys are used in aircrafts, as they are strong and light
Corrosion of metals
• Corrosion of metals is the attack of air, water or any surrounding substance to
the metals.
• Metals corrode when react with oxygen and other gases forming compounds.
Rusting:
- Only iron and steel can rust forming hydrated iron oxide Fe2O3.xH2O (flakey layer).
- Rusting is redox reaction.
2Fe2O3.2H2O(s)
4Fe(s) + 3O2(g) + 4H2O(l)
- Iron rusts faster in salty water (ionic compounds) as salts speed up oxidation
- Also high temperature speeds up rusting
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Prevent rusting:
1. Barrier method which is to cover the iron and keep it out of contact with
oxygen and water by painting greasing or covering with a different metal
(galvanization).
2. Let more active metal corrodes instead (sacrificial protection).
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Cathodic protection: The steel receives electrons from the battery, so it does not lose
its own electrons i.e. electrolysis process.
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1. The sacrificial protection of iron:
- Iron is used in in oil rigs and ships. It reacts with oxygen and water, forming Iron
(III) oxide or rusts.
- To prevent rusting, Iron is teamed up with a more reactive metal such as zinc. A
block of zinc may be welded to the side of ship.
- Zinc is more reactive than iron, loses electrons to the iron
2Zn(s)
2Zn2+ (aq) + 4e
(oxidation)
- Iron will pass electrons to oxygen and water vapour.
O2 (g) + 2H2O(l) + 4e4OH- The overall equation is
2Zn(s) + O2 (g) + 2H2O(l)
(reduction)
2Zn(OH)2 (aq)
- Zn is oxidized instead of iron. This is called sacrificial protection.
- Zinc block must be replaced before it all dissolves away.
2. Galvanization:
- This is another way to protect iron from rusting, where iron is coated with a layer of
zinc by electrolysis. The zinc coating keeps air and water vapour away. But if the
coating gets damaged, zinc will protect iron by sacrificial protection.
Recycling:
Recycling is the reuse of discarded materials after purification
Ex. Aluminium, copper, glass, papers and some plastics.
Recycling is very useful because:
1. It saves raw materials and our natural resources.
2. It decreases pollution.
3. It saves money needed to buy and extract ores especially in the electrolysis of
aluminium.
Disadvantages
• More transport on roads carrying used metals to recycling centers
• Energy consumed in collecting materials and sorting them per material type
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chloride
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Topic 9
Chemical Changes and Energy
Physical and chemical changes:
Physical change is a change in state or nature of the same substance.
Chemical change takes place during chemical reaction as new substance is formed.
Physical change
Chemical change
1-Easily reversed
- 1- Usually difficult to reverse
2-No new substance is formed - 2- New substance is formed
3- Energy changes usually
3-Energy changes usually
small and not significant
considerable and
significant
4. Usually called chemical
reaction.
Differences between mixtures and compounds:
Mixture
1- It contains two or more different
substances in any ratio.
2- No chemical change takes place
Compound
- 1- It is a single substance made of two
or more different elements
chemically bonded in a definite
ratio.
2. Involves a chemical change.
3- The components can be separated by
physical means
3. The components can be separated
by chemical means.
4- Keeps the properties of their
components.
4. Its properties are different from
those of its components.
5- No change in energy when the mixture
is formed.
5. Energy is given out or absorbed
6- Ex. Fe / S
6. Ex. FeS
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“Exothermic and endothermic reactions”
• Chemical reactions may be exothermic or endothermic.
A + B
Thermometer
AB
B
A
A
Ti
Tf
• During any chemical reaction, bonds of reactants are broken and bonds of products are
formed.
* Bond breaking
endothermic step
* Bond formation
exothermic step
“Exothermic Reaction”
Accompanied by release of energy [Tf > Ti].
A + B
AB + Heat
Hot
Surrounding
- The chemicals lose energy, the surrounding gain energy ∴Temperature increases
- Surrounding means reaction mixture, air in and around the beaker, the beaker itself and
the thermometer
Energy level diagram
N.B.
ΔH Enthalpy change means change in heat energy from reactant to product
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- Energy content of reactant is greater than energy content of product.
- The product is more stable than reactant.
Energy needed to break bonds of reactants is less than the energy produced
during formation of products.
Ex. Neutralization – Combustion – Displacement – Precipitation – Condensation –
Freezing
(any physical process takes place by cooling).
Activation energy is the minimum energy needed to start the reaction
“Endothermic Reaction”
Accompanied by absorption of heat [Tf < Ti].
A + B + Heat
AB
Cold
Surrounding
- The chemicals gain energy, the surrounding loses energy
∴Temperature decreases
Energy level diagram
- Energy content of reactant is less than energy content of product.
- The product is less stable than reactant.
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Energy needed to break bonds of reactants is more than the energy produced
during formation of products.
Ex. Decomposition [CaCO3
CaO + CO2]
– Dissolving of some salts [NH4Cl, KNO3] – Photosynthesis – Evaporation – Melting
(any physical process takes place by heating).
Starting the reaction off:
• For some reactions, not much energy is needed, the reaction takes place at room
temperature (spontaneous).
• Some exothermic reactions need heat from a Bunsen burner just to start bond breaking.
Then the energy given out by the reaction breaks further bonds.
• For endothermic reactions like the decomposition of calcium carbonate, you must
continue heating until the reaction is completed.
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Calculating Enthalpy change
H of a reaction:
Enthalpy is the heat content of the system
Enthalpy change is the transferee of thermal energy during a reaction
H = sum of bond energies of reactants – sum of bond energies of products
Ex.1 Calculate the heat of reaction and detect if the reaction is exothermic or
endothermic .
+
H2
Cl2
2 HCl
Given that the bond energies are :
H-H = 436 Kj
Cl-Cl = 242 Kj
H-Cl = 431 Kj
Answer:
H-H
436
+
Cl-Cl
242
2H-Cl
2X 431
Energy absorbed for bond breaking = 436 + 242 = + 678 Kj
Energy released for bond formation = 2 x 431 = - 862 Kj
Heat of reaction
H = 678 – 862 = - 184 Kj
H has a negative value
∴ the reaction is exothermic
Its energy level diagram is shown
Ex.2 In the following reaction
2 NH3
N2 +
3 H2
Calculate
H and find the type of reaction, if the bond energies are:
N = N = 946 Kj
N-H = 391 Kj
H-H = 436 Kj
Solution:
2 NH3
6x 391
N2 +
1x946
3H2
3x436
Energy absorbed for bond breaking = 6x391 = +2346 Kj
Energy released for bond formation = 946 + (3x436) = - 2254 Kj
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Heat of reaction H = 2346 - 2254 = +92 Kj
H has a positive value
∴ the reaction is endothermic
Ex.3
In the following reaction
C2 H 2 + H 2
C2 H 4
Calculate
H and find the type of reaction, if the bond energies are:
C = C = 838 Kj
C=C = 612 Kj
C-H = 412 Kj
H-H = 436 Kj
H H
Solution:
H-C=C-H
838
+
H-H
436
H- C= C-H
2x 412 + 612
Energy absorbed for bond breaking = 838 + 436 = + 1274 Kj
Energy released for bond formation = (2x 412) + 612 = - 1436 Kj
Heat of reaction H = 1274 – 1436 = - 162 Kj
H has a negative value
Ex. 4
∴ The reaction is exothermic
Show that the following reaction is exothermic.
CH4
+ F2
Bond energies are:
C-H = 413
F-F = 158
CH3F
+
H-F = 565
HF
C-F = 495 All in Kj/mol
Solution:
H
H- C-H
H
413
+
F-F
158
H
H-C-F
H
495
+
H-F
565
Energy absorbed for bond breaking = 413 + 158 = + 571Kj
Energy released for bond formation = 495 + 565 = - 1060 Kj
Heat of reaction H = 571 – 1060 = - 489 Kj
H has a negative value
∴ The reaction is exothermic
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Ex.4
Grade (10)
Calculate the heat of reaction and detect if the reaction is exothermic or
endothermic .
CH4
+
2O2
Given that the bond energies are :
C-H = 435
O=O = 497
Ex.5
CO2
C=O = 803
+
2H2O
H-O = 464
All in KJ /mol
Ethene burns in oxygen to form carbon dioxide and water vapour.
The bond energies are shown in the table.
What is the energy change for the reaction?
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Production of Heat Energy
• Fuel: is the substance that can be used as a source of energy.
Fuel
Fossil
-
Non-fossil
Ex: Coal
- Ex: Ethanol
Petroleum
- Hydrogen
Natural gas
Radio-active isotope U235(Which
Mainly made of C & H (except coal)
does not burn), used in nuclear
When burns CO2 + H2O
power stations
Non-renewable
- Some are renewable.
1. Non-renewable fuels:
(a) Fossil fuels:
• Fossil fuels are formed from the anaerobic decay of dead animals and plants under high
pressure and temperature. Plants turned to coal while animals are turned to oil.
• The incomplete combustion of fossil fuel produces carbon monoxide gas which causes
pollution.
• Fossil fuels contain impurities of sulfur which when burned, produced SO2 gas and
causes pollution (acid rain problem).
• Natural gas is mainly methane.
• Oil is preferred than coal as a fuel because oil is easier to be transported and stored and
can be used in different types of engines.
(b) Nuclear fuel:
• Uranium 235 is a radioactive isotope which is used as a fuel producing large amount of
heat through nuclear reactions. In a nuclear power stations, the heat formed turns water
to steam. The steam rotates a turbine connected to an electric generator to produce
electricity.
• It gives out huge amount of energy, without polluting gases as carbon monoxide
• Disadvantages: 1- waste material is a dangerous and radioactive, and very difficult to
find a place to store it safely.
2- An explosion may takes place, causing spreading of radioactive
material over a huge area, carried in the wind.
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2. Renewable sources of energy:
Ex. Solar energy, energy from wind, energy from water falls, ethanol& hydrogen
A good fuel must be:
1. Cheap.
2. Available in large quantities.
3. Safe to store and transport.
4. Easy to burn causing no pollution.
5. Release large amount of energy.
Ethanol as a fuel:
- Alcohol with the formula C2H5OH can be made from sugar cane or corn, and used as a
car fuel.
To compare between the heat energy released from burning two different fuels:
[N.B. The idea is: Heat given out by the fuel = Heat
gained by the water in the metal can]
1. Take a certain volume of water using a burette and
put it in a metal can.
2. Put a specific mass of fuel in a spirit burner.
3. Set the apparatus as shown below.
4. Use the thermometer to measure the initial
temperature of water T1.
5. Burn the fuel in excess oxygen
6. When the fuel is burned completely record the time, and the final temperature of water
T2.
7. Calculate, how much the temperature of water increased in a specific time
8. Repeat the experiment using the second fuel under same conditions [ water volume,
mass of fuel, Initial temperature of water, time of burning & distance between
the metal and the spirit burner]
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9. Compare the temperature rise for both fuels.
The fuel which gives a greater temperature rise in the certain time is the better one.
Possible errors:
1. Inaccurate volume of water.
2. The distance between the spirit burner and water can is not the same.
3. The thermometer may touch the metal can, so it would measure the temperature of
the can not the water.
4. Incomplete combustion due to shortage of oxygen
5. Mass of fuel is not the same.
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Hydrogen –Oxygen fuel cell
In electrochemical cells and batteries, Chemical energy changes to electric energy
[Exothermic]
Fuel cell is an electrochemical cell which is supplied continuously with oxygen and
hydrogen gases.
Hydrogen and oxygen combine without burning in a redox reaction. The energy is given
out as an electric current.
H2 (g) + 1/2 O2 (g)
H2O (i)
• Advantages and disadvantages of hydrogen cell are:
Advantages
1. Only water is formed, no pollution,
zero emission of carbon dioxide
2. Releases higher amount of energy
3. Works continuously i.e. no need for
recharging
4. Renewable
5. Virtual emission free
6. Nontoxic
Disadvantages
1. Large fuel tank required
2. Few filling car stations
3. Special car engines needed to suits
fuel cells
4. Expensive
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Topic 10
Rate of Reaction
• Chemical reaction takes place when: reacting particles collide with enough energy
to form the product (successful collision).
• Activation energy: is the energy needed to activate the reactants, break the bonds
and start the reaction.
• Some reactions are very fast, others are slow.
Examples of fast reactions:
1- Explosions and fireworks.
2- Precipitation reactions as solutions of ionic compounds react quickly
3- Coal mining
Examples of slow reactions:
1- Rusting of iron.
2- Formation of fossil fuel.
A + B
Reactants
C + D
Products
• Reaction rate: is a measure of how fast the reaction takes place.
Rate =
Change in concentration of reactant or product
Time
[The rate of reaction is inversely proportional to the time]
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To find the rate of reaction, we should measure:
- Concentration of reactant used up per unit of time or
- Concentration of product produced per unit of time
♣ How to measure the rate of a reaction?
•
•
•
•
•
Volume of gas produced.
Decrease in mass of reagent due to a gas release.
Colour change.
PH change.
Temperature.
(1) Reactions give a gas as one of the products:
a) Measuring the volume of gas produced / unit time:
Mg(s) + 2HCl(aq)
MgCl2(aq) + H2(g)
Magnesium
N.B. HCl is excess which means” more than enough to react”.
- Test tube containing magnesium to separate the reactant, adjust the start of reaction
and avoid timing error.
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min.
Total volume of hydrogen
The average rate of reaction = --------------------------------Total time for the reaction
N.B. The volume of hydrogen is:
12 cm3 after the first minute.
24 cm3 after the second minute.
12 cm3 in the second minute.
b) Measuring the mass loss / unit time:
CaCO3 + 2HCl
CaCl2 + H2O + CO2
N.B. Balance reading decreases as CO2 evolves.
Cotton wool plugs to prevent splashing of liquids and allows gas to escape.
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Mass
(gm)
100 –
50 –
‫׀‬
1
0
‫׀‬
2
‫׀‬
3
‫׀‬
4
‫׀‬
5
‫׀‬
6
Time (min.)
Total mass loss
Average rate of reaction = ------------------------------Total time of reaction
(2) Reactions give precipitate
Na2S2O3 + 2HCl
2NaCl + SO2 + H2O + S
Sodium thiosulfate
Yellow
The cross disappears when enough sulfur is has formed to hide it’
• This experiment is used to study the effect of changing temperature or concentration
on the time of disappearance of the cross and consequently the rate of reaction
• Each time the total volume of reactants, the conical flask and the printed paper should
be the same
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Factors Affecting the Rate of Chemical Reaction
Sometimes, it is useful to increase or decrease the rate of reaction so we have to study
the factors affecting rate of reaction.
i.e. Increase the rate of reaction is due to increase the chance of successful collision
1) Temperature:
When temperature increases: the particles gain thermal energy; more particles have
energy above activation energy and move faster. Number of successful collisions per
time increases. The collisions are stronger
∴ Rate of reaction increases
N.B. The low temperature in the fridge slows down reactions that make food rote.
2) Particle size:
[solids]
- By decreasing the particle size (increasing surface area): more reactant particles
exposed to collide, and more frequent successful collisions increase.
∴ Rate of reaction increases
Coarse or Lumps
Small granules
“Rate increases”
Powder
“Time decreases”
Mass
V
Powder
lumps
lumps
T
T
Powder
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Explosion:
Explosion is a dangerously chemical reaction.
- In flour mills: large surface area of flammable flour dust
catches fire easily; a spark from a machine could be
enough to cause an explosion.
For the same reason, explosions are a risk in wood mills,
from wood dust, and in silos where wheat and other
grains are stored. And in factories that make custard powder, and dried milk. The dust
from all these will burn.
- In coal mines: when methane and other flammable gases reach certain concentration,
they form an explosive mix with air. A spark is enough to set off an explosion.
3) Pressure:
[gases]
- When the pressure increases: the particles come closer, more particles in same
volume and more frequent successful collisions take place.
∴ Rate of reaction increases
3H2 + N2
2NH3
Total volume
4 moles
2 moles
• When pressure increases, the reaction tends to the direction of formation of NH3 ,
i.e. less number of moles (less volume).
• When pressure decreases, the reaction tends to the direction of formation of H2 & N2
i.e. high number of moles (more volume).
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4) Catalyst:
• Catalyst: is a substance that speeds up the chemical reaction without being
chemically changed.
- By adding a catalyst: it lowers the activation energy, more particles have enough
energy to successfully collide.
∴ Rate of reaction increases
Without
catalyst
E
V
Mass
Catalyst
A+B
With
catalyst
Catalyst
AB
T
T
MnO2
Example:
2H2O2
2H2O + O2
- Hydrogen peroxide H2O2 is a colourless liquid that breaks down very slowly to
water and oxygen.
- Manganese (IV) oxide is used as a catalyst to speed up the reaction thousands
of time.
- The more catalyst is added, the faster the reaction goes
N.B.
Enzymes
are biological catalysts, protein in nature; work at limited range of
temperature & pH, increasing rate of reaction.
Amylase in mouth speeds up the breakdown of starch
Biological detergents contain enzymes that help to break down grease, food stains
and blood stains on clothing in the wash.
5) Concentration:
- By increasing concentration, more particles present in same volume to react, more
frequent successful collisions take place. ∴ Rate of reaction increases
Lower concentration
higher concentration
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d
X
• Generally increasing concentration, increases number of particles in same volume,
and more frequent successful collision take place so rate increases.
• [a curve] Increasing concentration of excess reactant, increases rate of reaction only
which means steeper curve at start while volume of gas produced will be the same.
• [d curve] Increasing concentration of limiting reactant, increases both rate of reaction
and volume of gas produced.
• [b curve] Increasing volume of limiting reactant, increases volume of gas produced
but not the rate of reaction.
• [c curve] Halfling concentration of limiting reagent, decreases rate of reaction which
means less steep curve and reduces volume of gas produced to half.
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Reversible Reactions
A + B
C + D
• Reversible reaction: can proceed in both directions forward and backward.
Example (1) :
Water vapour
Heat
CuSO4.5H2O
CuSO4 + 5H2O
Hydrated
copper sulfate
(blue)
Anhydrous
copper sulfate
(white)
N.B. Endothermic reaction.
CuSO4 + 5H2O
(white)
Crystallization
Cool
CuSO4.5H2O
(blue)
N.B. Exothermic reaction.
CuSO4.5H2O
Cool
CuSO4 + 5H2O
N.B. Very important test for presence of water.
Example (2)
3H2 + N2
2NH3
Chemical equilibrium: is the state when rates of forward and backward reactions are
equal, where the concentrations of reactants and products are constant.
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Le Chatelier’s Principle:
When a change is made to the conditions of a system in dynamic equilibrium, the system
moves to the direction that opposes the change and retains equilibrium.
Factors affecting a reaction at equilibrium:
1. Concentration.
2. Pressure.
3. Temperature.
1. Temperature:
a] Exothermic reactions:
A + B
C + D + Heat
∆ H = - ve
-By increasing temperature, the equilibrium shifts backwards, concentration of
A& B increases; C& D [yield] decreases.
-By decreasing temperature [cooling], the equilibrium shifts forwards,
Concentration of A& B decreases; C& D [yield] increases.
b] Endothermic reactions:
A + B + Heat
C + D
∆H = + ve
-By increasing temperature, the equilibrium shifts forwards, concentration of
A& B decreases; C& D [yield] increases.
-By decreasing temperature [cooling], the equilibrium shifts backwards,
concentration of A& B increases; C& D [yield] decreases.
2. Pressure:
2SO2 +
[for gases]
O2
2 SO3
-By increasing pressure, the equilibrium shifts forwards[less number of moles],
concentration of SO2& O2 decreases, SO3 [yield] increases.
-By decreasing pressure, the equilibrium shifts backwards [more
number of moles] concentration of SO2& O2 increases, SO3 [yield] decreases.
N.B. If number of moles of reactants and products are equal, the pressure has no
effect.
N2
+ O2
2NO
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3. Concentration:
[for liquids]
A + B
C + D
-By increasing concentration of A or B, the equilibrium shifts forwards,
and concentration of C& D [yield] increases.
-By decreasing concentration of A or B, the equilibrium shifts backwards,
and concentration of C& D [yield] decreases.
Ex.
BiCl3 (aq) + H2O (l)
Colourless
BiOCl(s) + 2HCl(aq)
White
Using catalyst: does not affect the position of equilibrium, it increases both
forward and backward reactions by lowering the activation energy, and only the
reaction reaches the equilibrium faster which saves time.
Haber process:
N2 (g)
+
3H2(g)
Forward reaction is exothermic
2NH3(g) + Heat
∆H = - ve value
Conditions of reactions are:
• Iron powder catalyst.
• 450 °C temperature.
(Optimum temperature, low enough to get a reasonable yield but high enough to
get a fast reaction rate and be economical)
• 20000 kpa / 200 atm. Pressure.
(Optimum pressure, high enough to get a reasonable yield but low enough to
be economical)
• Nitrogen& hydrogen are mixed in a ratio 1 : 3 by volume.
• Source of nitrogen is fractional distillation of liquid air
• source of hydrogen is
1. Electrolysis of brine where hydrogen is collected at cathode
2. Reaction of methane with steam in presence of nickel as catalyst
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CH4 (g) + H2O(g)
CO(g) + 3H2(g)
Carbon monoxide formed could poison the catalyst, and therefore removed by
reaction with more steam
CO(g) + H2O(g)
CO2(g) + H2(g)
3. BY cracking of hydrocarbon
• The mixture of the two gases will never react completely; the yield will never be
100%. Hydrogen and nitrogen are sent to be recycled and react once more.
Ammonia
- Ammonia is the world’s second most manufactured chemical after sulfuric acid.
- Can be prepared in lab by heating any ammonium compound with a strong base
2NH4Cl(s) + Ca(OH)2(s)
CaCl2(s) + 2H2O(l) 2NH3(g)
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- The properties of ammonia:
i. Colourless gas with a strong, choking smell
ii. Less dens than air.
iii. Reacts with dilute hydrogen chloride gas to form a white smoke. This
reaction can be used to test for presence of ammonia.
iv. Very soluble in water forming alkaline solution.
- Ammonia is used to make fertilizers and nitric acid.
Fertilizers: are substances added to the soil to make it more fertile
Animal manure is a natural fertilizer. Synthetic fertilizers are made in factories
Synthetic fertilizers are minerals which are added to the soil to promote plant growth
and contain Nitrogen, phosphorous and potassium[N , P , K]
Ex. Ammonium phosphate, Potassium nitrate
♣ Fertilizers are used to:
1. Get better growth of plant.
2. Get better yield.
3. Compensate used nutrients in the soil.
N
To make chlorophyll and proteins
P
To help leaves and roots growth and crops to ripen.
K
To promote growth and resist diseases
It’s not all good news. There are drawbacks for using fertilizers:
1. Fertilizers seep into rivers from farmland, help algae to grow which when die
, oxygen decreases in water and bacteria that feed on them increases so fish
suffocate
2. Nitrate ions from fertilizers can end up in our water supply, changing to
nitrite ions that react with hemoglobin instead of oxygen casing illness
spatially in infants.
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Contact process
To prepare sulfuric acid, sulfur is roasted in oxygen
S(s)
+ O2 (g)
SO2 (g)
2SO2 (g) + O2 (g)
2SO3 (g)
= - ve
Sulfur dioxide is further oxidized to form sulfur trioxide which is called contact
process. The reaction is exothermic
Conditions of reactions:
1. Temperature = 450°C
(Optimum temperature, low enough to get a reasonable yield but high enough to
get a fast reaction rate and be economical)
2. Pressure = 200 kpa / 2 atm.
( Although theoretically increasing pressure, increases the yield as the equilibrium
position shifts to the RHS which has less number of moles, practically the increase
in yield will be limited i.e. economically does not worth. Fair enough to use Catalyst
and relatively high temperature.)
3. Vanadium V oxide catalyst.
SO3 (g) + H2SO4(aq)
H2S2O7 (l)
Oleum
H2S2O7(l) + H2O(l)
2H2SO4 (aq)
N.B. Sulfur trioxide is passed over concentrated sulfuric acid [98%] forming oleum
[99.5%]
Sulfur trioxide is not passed over water as it is a highly violent exothermic reaction
forming a toxic cloud.
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Topic 11
Chemistry of environment
Environment is the layer of air and water that surrounds the earth
Air quality and climate
78.1 % N2
20.95% O2
0.04% CO2
H2 O
0.3% Inert gases (Ar)
­ Components of air can be separated by fractional distillation of liquid air.
­ Liquid air is warmed up and the gases boil at different temperatures and then collected
one by one.
Nitrogen is collected first (b.p. ­196°C) Followed by Argon (b.p.­186°C), then oxygen
(b.p.­ 183 °C).
• Air pollution: Is any change in air components causing harms for living organisms.
• Air pollutants: Carbon dioxide / carbon monoxide / Particulates / Oxides of Sulfur /
Oxides of Nitrogen / Methane
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Pollutant
Source
Carbon dioxide
Colourless,
odorless and
tasteless
• Carbon monoxide:
Colourless,
odorless and
insoluble
Particulates
( soot)
Complete combustion
of carbon­based fuels
Sulfur dioxide:
Acidic gas with
sharp smell
♣
Oxides of
nitrogen
Methane
Adverse
effects
Action to
reduce the
pollutant
Global warming and
climate change
Reduce use of
carbon­based fuels
and increase green
areas
Use a catalytic
converter
Incomplete
Health hazard: prevents
combustion of carbon­ blood from transporting
based fuels
oxygen
Very tiny particles
produced during the
incomplete
combustion of fuel
Burning fossil fuels
that contain sulfur
mainly in power
stations.
Oxygen and nitrogen,
which react at high
temperatures in car
engines
From vegetation and
waste gases from
digestion of animals
Use of fine mesh
filters in diesel
vehicles.
­Acid rain, causing
deforestation, damage to
buildings
­respiratory problems
Health hazard: causes
severe respiratory
problems; photochemical
smog and acid rain
Global warming and
climate change
Use low sulfur
fuels a
desulfurization
Use a catalytic
converter
Reduces rice
farming
• Carbon monoxide is formed due to incomplete combustion of petroleum fuel
2CO(g) + 4H2O(l)
2CH4(g) + 3O2(g)
Carbon particulates are formed due to incomplete combustion of carbon based fuel
partially in diesel vehicles.
2C8H18(g) + 9O2(g)
16C(s) + 18H2O(l)
Octane
Oxygen
Particulate (Soot)
Water
Acidic rain is rain which is more acidic then normal due to presence of dissolved
pollutants as sulfur dioxide and nitrogen dioxide.
Desulfurization is the removal of sulfur dioxide (acidic gas) from the flue (chimney)
by neutralization with calcium oxide (base).
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Grade (10)
Oxides of nitrogen
­ They are formed due to reaction of air components [Oxygen& nitrogen] at high
temperature inside the car engine.
­ Oxides of nitrogen dissolved in water forming nitric acid
♣ Photochemical smog results when several environment pollutants like nitrogen dioxide
react with sunlight form brown gas seen over many large cities.
It has been linked to respiratory problems and asthma attack.
∗ Lead compounds are added to fuel to help it burns smoothly in engines.
Carbon monoxide, oxides of nitrogen and unburned hydrocarbons are the car
exhaust
Reducing air pollutants:
1. In modern power stations slaked lime (basic) is used to treat acidic sulfur dioxide
2. Most countries have now banned lead in petrol. But it can still arise from battery
factories.
3. The exhaust of new cars is fitted with catalytic converter, in which harmful gases are
converted to harmless gases.
4. Gas heaters and boilers must be checked regularly, to make sure that air supply is not
blocked by soot.
• Harmful CO& oxides of nitrogen react together inside the catalytic converter to form
Harmless CO2 & N2 gases.
2NO + 2CO
N2 + 2CO2
• The catalytic converter is made of a ceramic honeycomb to provide a large surface area,
coated with a layer of palladium and rhodium as a catalyst.
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Carbon dioxide
Carbon dioxide is a greenhouse gas, produced naturally in air as a product of three
reactions:
1. Complete combustion of carbon based fuel:
Burning carbon based fuel releases carbon dioxide
CH4 (g) + 2O2 (g)
CO2 (g) + 2H2O (l) + Energy
2. Respiration:
Takes place in living cells to provide energy. This energy keeps us warm and very
important for all biological processes.
C6H12O6 (aq) + 6O2 (g)
6CO2 (g) + 6H2O (l)
Carbon dioxide is removed from air by photosynthesis:
Carbon dioxide and water react in plant leaves, to give glucose and oxygen.
Chlorophyll, a green pigment in leaves, traps sunlight and acts as a catalyst for the
reaction,
6CO2 (g) + 6H2O (l)
Carbon dioxide
water
Light
Chlorophyll
C6H12O6 (s) + 6O2 (g)
glucose
Oxygen
The plant uses the glucose to make the other compounds it needs .The animals eat the
plans. So carbon compounds get passed along the food chain to humans.
Properties of carbon dioxide:
1. Colourless and odorless gas
2. Much heavier than air
3. Extinguish fires
4. Slightly soluble in water forming carbonic acid(H2CO3)
CO2 is a greenhouse gas but not a poisonous one.
Greenhouse gases, and global warming
• Greenhouse gases absorb reflected thermal energy from the earth, and prevent it from
escaping into the atmosphere, which reduces heat loss to space and increases the
temperature of lower atmosphere
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• There are several greenhouse gases water vapour, oxides of sulfur and nitrogen and
chlorofluorocarbon (CFCs) are also greenhouse gases, but carbon dioxide and methane
are the two main ones
• The level of carbon dioxide in atmosphere is rising because we burn more fossil fuel
each year.
• The level of methane is rising because of
1. The increase in animal farming as methane is a part of their digestion wastes.
2. More landfill sites where waste food decomposed anaerobically by bacteria
producing large amount of methane.
3. Rice farming.
• Advantage of greenhouse gases is providing a suitable temperature at night and protect
us from freezing in absence of sunlight
• Greenhouse effect: is a natural phenomenon in which, thermal energy reflected from
the surface of the earth is trapped by greenhouse gases keeping the temperature of the
earth relatively constant
• Global warming: is the increase in average temperatures around the world due to the
increase in percentage of greenhouse gases
• Climate change: air temperature affects rainfall, and cloud cover, and wind patterns.
The scientists predict that:
­ Areas with heavy rainwater will become very dry, which may increases the
probability of crop failure. Other places will get much wetter.
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­ Melting of earth’s polar ice caps will cause sea levels to rise, so low­lying countries
will be at risk of flooding, and faster rate of coastal erosion
­ Species that cannot adapt to the changing climate will die out ex. Polar bears
­ Also migratory patterns of animals and birds will change.
­ Storms, floods, landslides and wildfires will be more frequent and severe.
­ Changing in sea temperature can lead to bleaching coral reefs and loss of marine life.
­ Disturbance of life cycle of some living organisms who are sensitive to any change
in temperature.
• What can we do?
We have to reduce new emission of carbon dioxide, to stop warming getting out of
control. There have been two important climate change conferences Kyoto Protocol of
2005 and the Paris agreement of 2016
­ Reduction of our reliance on fossil fuel for transportation and electricity generation.
­ People should use public transport or bikes or walking rather than going by car.
­ Countries should use clean ways to get electricity like wind power and solar power
from renewable resources.
­ Increase the interest in moving away from petrol cars to electric cars and hydrogen
fuel cell vehicles
­ At generating electricity plants, it may be possible to capture carbon dioxide before it
is released to the atmosphere, and store it under ground.
­ Scientists should find new ways to reduce amount of CO2 entering the atmosphere. For
example capturing it from power stations and chimneys.
­ Planting additional trees to capture CO2 for photosynthesis.
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Water
• 70 % of Earth’s surface is covered with sea water, beside lakes and rivers.
• 72 % of human bones is water
• 82 % of human kidneys is water
• 90 % of blood is water
• Pure water is a neutral, colourless liquid at 1 atm., boils at 100 Cº and freezes at 0 Cº, and
used in experiments, as impurities
­ may interfere with experiments giving unwanted side reactions
­ may interfere with the results in standard analysis
• Tap water in not a pure water, as it contains dissolved salts.
• The unique properties of water:
1. Excellent solvent for ionic compounds
2. It has high boiling point in spite of having low relative molecular mass.
3. It has high specific heat capacity
4. It decreases in density when freezes.
Chemical test for presence of water:
Anhydrous copper sulfate changed from white to blue hydrated copper sulfate.
CuSO4 + 5H2O
CuSO4.5H2O
White
Blue
Anhydrous cobalt chloride changed from blue to pink hydrated cobalt chloride.
CoCl2 + 6H2O
CoCl2.6H2O
Blue
Pink
Test for purity of water :
Measure the melting point or boiling point .It will be sharp if pure.
Sources of water: rivers and ground water. Water contains mud particles, animal
wastes, bits of dead vegetation and microbes like bacteria.
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Uses of water:
Industrial uses
Domestic uses
1. Washing, cleaning raw materials.
2. Food processing.
3. Cooling.
4. Electric power stations (the steam drives
the turbines that generate electricity).
1. Drinking and cooking.
2. Washing and cleaning.
3. Flushing toilets.
4. Drink for animals and water crops
Water pollution and treatment
Water from natural resources contain beneficial and harmful substances.
Beneficial substances:
• Dissolved oxygen and carbon dioxide
­ Oxygen is very important to support plant and animal life
­ Oxygen enters the water as a result of photosynthesis of aquatic plants or by diffusion
from air.
­ Oxygen is removed by respiration of aquatic plants and animals
­ Carbon dioxide creates low level of acidity in water, increased by vehicles and
industry
• Metallic compounds (minerals)
­ Minerals are dissolved metallic compounds from the rocks that are needed in small
amounts for good health.
­ These include group I metal ions [sodium and potassium], group II metal ions
[calcium and magnesium], and transition metal ions[Fe, Co, Cu, Ni, Zn, and Cr]
­ Calcium supports the health of teeth and bones, while iron is needed in production of
hemoglobin.
Harmful substances:
• Some harmful metallic compounds
­ Heavy toxic metals [Pb, Cd and Hg] may enter water system via mining, metal
smelting, waste disposal, corrosion and metal Processing plants.
­ Lead can cause liver and kidney damage while mercury affects the nervous system
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• Sewage:
­ Human waste water usually treated in certain plants to remove harmful materials.
­ Leaks of sewage into drinking water during natural disaster like earthquakes or severe
weather events.
­ Harmful microbes enter drinking water spreading diseases such as diarrhea, cholera,
dysentery, typhoid and polio.
• Nitrates and phosphates:
­ Run off water containing NPK fertilizers, some pesticides and phosphate detergents
used in home and industry may washed over the surface of the soil into the rivers
­ They cause the rapid growth of algae forming huge blooms that cover the surface of
water that turned green and block out sunlight. Aquatic plants cannot make
photosynthesis and will die [eutrophication].
­ Shortage of oxygen level causes death of many aquatic animals
• Plastics
They are poor disposal due to their large quantity, and being nonbiodegradable, these
cause many problems:
1. Fish, whales, turtles, and other animals eat them, damaging their digestive system and
starve to death.
2. Large sea creatures and sea birds may be trapped by discarded fishing nets.
Purification of domestic water
This takes place in four steps:
1. Screening is the removal of large insoluble solid objects like rocks and
plastic bags by sedimentation
2. Aluminium sulfate is added to coagulate small particles of clay to form large
clumps which settle then, filter to remove solid insoluble substances
3. Use of carbon to remove tastes and odors.
4. Chlorination to kill microbes and harmful bacteria.
5. Fluoride is sometimes added to decrease tooth decay
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TOPIC 12
Organic Compounds
Organic chemistry is the study of properties and chemical reactions of carbon
compounds.
­ Carbon is tetravalent.
–C–
C=
­ Carbon atoms may link together to form a chain.
–C≡
–C–C–C–
Organic compounds are:
1. Hydrocarbons Organic compounds made of C & H only
2. Hydrocarbon derivatives made of C, H & S or O or N
3. Carbon is unique in the variety of molecules it can form as
Carbon atoms can join to each other to form long chains
The carbon atoms in a chain can be linked by single, double or triple covalent bonds
Carbon atoms can also be arrange themselves in rings
“Fuels”
Substances which are used to release heat energy.
• Coal, oil and natural gas are fossil fuel.
• Coal has been formed by the anaerobic decay of vegetation over millions of years.
• Petroleum or crude oil is a smelly mixture of hundreds of different hydrocarbons. In
petroleum, hydrocarbon molecules have different shapes and sizes, with different
numbers of carbon atoms, from 1 to 70.
• Petroleum is a non­renewable resource. It is expecting that world’s reserve will last
about 40 years
• Refining oil is the separation of oil into groups of compounds having similar size
boiling point.
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Fractional distillation:
• Fractional distillation separates petroleum into more useful mixtures of hydrocarbon.
• Fraction is the distillate collected over a narrow range of temperature from the
fractionating tower.
• As we go down the tower:
­ Boiling point increases
­ less volatile
­ Relative molecular mass increases
­ less flammable
­ Chain length increases
­ More viscous (thick & sticky)
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N.B. As we go down the tower boiling point of distillates increases as the chain length
increases and intermolecular attraction forces increase.
The fractions all need further treatment before they can be used:
­ Sulfur impurities should be removed
­ Some fractions are separated further into single compounds
­ Part of fractions may be cracked (breaking molecules down into smaller ones)
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“Naming Organic Compounds”
• The prefix of the name tells the length of the chain (ALK).
• The suffix of the name tells the type of covalent bond between carbon atoms.
Number of
C
ALK
ALKANE
Meth
│
Methane ─ C ─
│
Eth
│ │
Ethane ─ C ─ C ─
│ │
3
Prop
Propane
│ │ │
─C─C─ C ─
│ │ │
4
But
Butane
Butene
5
Pent
Pentane
Pentene
6
Hex
Hexane
Hexene
7
Hept
Heptane
Heptene
8
Oct
Octane
Octene
1
2
ALKENE
Ethene
C═C
Propene
│ │
─ C ─ C ═C
│
♣ The organic compounds can be represented by two ways:
Molecular formula: represents number and type of atoms forming one molecule.
Displayed formula: represents the arrangement of atoms and type of bonds in one
molecule. The bonds are represented as lines.
Structural formula: all atoms are indicated using subscript numbers, not all bonds
are shown. Carbon hydrogen bonds are often simplified.
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“Alkanes”
Alkane is a homologous series of saturated hydrocarbon with a general formula
(CnH2n+2).
•
•
•
•
Alkanes have only C to C single bond.
General formula [CnH2n+2].
Saturated hydrocarbon.
Generally unreactive [strong bonds] except for combustion and chlorination.
Physical properties:
1. The first 4 members are gases, the next 12 are liquids, and the rest are solids.
2. Boiling point increases with chain length because attraction forces between the
molecules increase.
Chemical properties:
(1) Combustion:
­ Reaction with excess O2 to produce CO2 and H2O and heat.
• N.B. Burning is exactly like combustion but when flame is developed.
CH4(g) + 2O2(g)
CO2(g) + 2H2O(g)
2C2H6(g) + 7O2(g)
4CO2(g) + 6H2O(g)
­ The incomplete combustion of alkane produces carbon monoxide and water.
(2) Substitution:
H
H
H ─ C ─ H + Cl ─ Cl
H
Methane (g)
Light
H ─ C ─ Cl + HCl
H
Chlorine (g)
Chloromethane (g) Hydrogen Chloride(g)
• It is also a photochemical reaction where activation energy needed to start the reaction
is provided by light
• Substitution reaction: Atom (or atoms) of a molecule is (are) replaced by different
atom(s) without changing the general structure of the molecule.
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♣ Uses of alkane:
1. Making of
2. Preparation of
3. Fuel, Wax and lubricating
alkenes.
hydrogen.
agent (oil)
♣ Structural isomerism:
­ Different organic compounds having same molecular formula but different structural
formula.
Ex. C4H10
CH
CH
3 3
─C─C─C─C─
&
─C─C─C─
C4H8
─C=C─C─C─
─C─C=C─C─
“Alkenes”
Alkenes are homologous series of unsaturated hydrocarbon with a general formula
(CnH2n).
•
•
•
•
The carbon chain has at least one double bond (C=C), unsaturated hydrocarbon.
General formula [CnH2n].
Highly reactive [addition reaction takes place].
First 3 members are gases at room temperature.
How to prepare alkenes?
By cracking of alkanes.
Cracking:
­ Is the breaking down of a long chain hydrocarbon to form smaller ones.
Condition of reaction is heat and catalyst.
SiO2 + Al2O3
C10H22 (g)
Decane (naphtha)
C8H18 (g) + C2H4 (g)
Octane
Ethene
C10H22 (g)
C7H16 (g) + C3H6 (g)
Heptane
Propene
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C10H22 (g)
♣ Cracking used to make:
­ Alkene to form polymers.
C7H14 (g) + C3H6 (g) + H2 (g)
Heptene
Propene
­ Hydrogen
­ Petrol(octane)
(1) Addition reaction: Forms only one product.
Turns an unsaturated alkene into a saturated compound.
♣ Hydrogenation Reaction
H
H
H
H
H2
C═C
H ─ C ─ C ─H
Ni / 180°C
H
H
H H
Ethane(g)
Ethene(g)
♣ Catalytic hydration of ethene:
H
H
H OH
300°C / 60 atm
C═C +
H ─ OH
H─C─C─H
Phosphoric acid
H
H
H H
Ethene (g)
Ethanol(g)
♣ Test for alkene:
­ Add few drops of bromine water, the colour changes from brown to colourless.
H
H
C═C
H
Br Br
+
Br ─ Br
H
Ethene (g)
Bromine (aq)
H─C─C─H
H H
1, 2 dibromoethane
(2) Addition polymerization:
­ Polymerization is the formation of long chain molecules
H
H
H H
n
C═C
─C─C─
n
H
H
H H
Ethene
Polyethene
[C2H2n]
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­ Alkenes are used: In making polymers and alcohol.
“Alcohols”
Alcohol is a homologous series of hydrocarbon derivatives of general formula
(CnH2n+1OH).
• OH is the function group.
• Alkanol and the number in the name give the position of OH.
OH
CH3CHCH3
[2-propanol]
CH3CH2CH2OH
[1-propanol]
Preparation of alcohols:
(1) Catalytic hydration of ethene (the chemical way):
O─H
C═C
Ethene (g)
+
H ─ OH
Water(g)
300°C / 60 atm
Phosphoric acid
─C─C─
Ethanol(g)
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The reaction is reversible and exothermic, high pressure and low temperature would
give the best yield. But in practice the reaction is carried out under optimum conditions
300°C, 60 atm. and catalyst to give a decent rate.
(2) By fermentation of sugary solution (the biological way):
­ Anaerobic fermentation by yeast fungus at 37°C and absence of O2.
­ 37°C if higher temperature, the enzymes will be denatured.
­ No oxygen to avoid oxidation of ethanol to ethanoic acid.
­ The reaction stops after certain time because the % of ethanol reached the level that
killed the yeast or the mixture gets too warm.
­ The mixture is filtered to remove yeast then fractional distillation takes place to
separate the alcohol.
C6H12O6(aq)
Yeast fungus
Fermentation
Ethanol by fermentation
Advantages
2C2H5OH (aq) + 2CO2 (g) + energy
Ethanol by hydration
• Use of renewable resources
• Fast reaction
• Waste plant material can be used • Run continually, just by removing
ethanol
• Produces pure ethanol. No need for
fractional distillation
Disadvantages • A lot of plant material used to • Made
from
non­renewable
make a liter of ethanol
resources.
• Slow reaction.
• Heat is needed which casts money.
• Yeast stops working after certain • The reaction is reversible. So the
time – even if glucose is left. Heat unreacted ethene and water must
of fraction distillation costs keep recycling.
money
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­ Alcohol is flammable, burns with clean blue flame
2CO2(g) + 3H2O(g)
C2H5OH (l) + 3O2(g)
­ Less pollutant than petrol as no SO2.
­ Flammable, burns with clean blue flame
♣ Ethanol is used:
­ As fuel for cars (biofuel), as it is quite cheap, made from waste plant material
­ Solvent in perfumes, glues (volatile) and food industry.
­ Alcoholic drinks.
­ Antiseptic.
“
Carboxylic Acids”
Organic or carboxylic acid is a homologous series of hydrocarbon derivatives of
general formula (CnH2n+1 COOH).
• Weak acid partially ionized
[pH ≈ 6]
• The function group is the carboxylic group (─ COOH)
• The name is alkanoic acid.
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Ethanoic acid: CH3COOH
H O
H─C─ C─O H
H
• Ethanoic acid can be prepared by oxidizing ethanol.
[O]
C2H5OH
CH3COOH
• The oxidation can be carried out in two ways.
Preparation of Ethanoic acid:
1) By fermentation- biological way:
Ethanol is left standing in air; bacteria bring about its oxidation to Vinegar or
Ethanoic acid (acid fermentation).
2) Using an oxidizing agent- Chemical way:
• Ethanol is oxidized much faster by warming it with the powerful
oxidizing agent acidified potassium manganate (VII). The
manganate (VII) ions Re reduced to Mn2+ ions, with colour
change. The acid provides the H+ ions for the reaction.
• Reflux technique is used to condense vapours.
MnO4- + 8H+ + 5e­
Purple
Mn2+ + 4H2O
colourless
Chemical properties:
1. Displacement reaction reacts with metal forming salt and hydrogen
Mg(s) + 2CH3COOH (aq)
Ethanoic acid
(CH3COO)2Mg (aq) + H2(g)
Magnesium ethanoate
2. Neutralization reaction
NaOH (aq) + CH3COOH (aq)
CH3COONa (aq) + H2O(l)
Sodium Ethanoate
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3. Esterification is a reaction between carboxylic acid and alcohol
Acid (l) +
HO
H­C­C­OH +
H
Ethanoic acid(l)
Alcohol(l)
Ester(l) + Water(l)
H
HO­C­H
HO
H­C­C­O­C­H
H
Methanol(l)
H
H
+
H2O
H
Methyl ethanoate(l)
Water(l)
• Esterification is a condensation reaction
• The reaction is reversible; sulfuric acid acts as a catalyst.
• The alcohol part comes first in the name­ but second in the formula (methyl ethanoate
/ magnesium ethanoate).
• Many esters are having attractive smells and tastes. So they are added to shampoos
and soaps, and ice cream and foods as flavorings.
• Esters: is a family of organic compounds formed due to esterification reaction and
have pleasant taste and smell.
• The name of ester is alkyl alkanoate. [Alkyl is from the alcohol and alkanoate is
from the acid]
Homologous Series:
♣ Is a family of similar compounds having:
1. Same functional group.
2. Similar chemical properties and reactions.
3. General formula.
4. Made by similar chemical reactions.
5. Their physical properties are predictable.
6. Each member differs by CH2 from the next one.
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Polymers
• Macromolecule: is a very large molecule made of repeating units.
• Polymer: is a very large molecule made of many monomer molecules.
• Monomer: is a small molecule that join together forming polymer.
• Addition polymerization: is the formation of polymer only by breaking double bond
of monomers. The monomers must have double bonds.
Cl
n
H
C=C
H
Cl H
Δ / pressure
Catalyst
H
Chloroethene
—C–C—
H H n
poly chloroethene (PVC)
Plastic
The chains are not all the same length. This is why we cannot write exact formula for
PVC or any polymer.
Plastics:
­ Are a group of polymeric materials characterized by their elasticity, ability to be
molded and shaped under heat and pressure.
­ Most plastics are made from chemicals in the naphtha fraction of petroleum
­ Can be molded into shapes without breaking.
(PVC): Used in hoses, water pipes & electric insulators.
Polyethene: Can be used in making bowls, plastic bottles and plastic bags.
• Condensation polymerization: a reaction in which there are two products, the
polymer and a smaller one.
­ Does not depend on C=C bonds
­ Two types of monomer join. Each has two function groups.
­ They join at their function groups by eliminating a small molecule.
Example: in making synthetic fibers
(Teryelne & Nylon).
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Advantages of synthetic fibers and plastics:
1. Do not conduct heat or electricity(insulators)
2. Flexible and can be coloured.
3. Do not corrode, and durable (not affected by air water and chemicals) water proof.
4. Strong because their molecules are attracted to each other, but also low density
5. Their properties can be changed by changing the monomers and the reaction
conditions.
Disadvantages of synthetic fibers:
Polyethene is the biggest problem. It is the most used plastic in the world as 5 trillion
plastic bags are made every year. So they cause many problems:
1. Fish, birds, and other animals eat them and starve to death.
2. They clog up drains and sewers, and cause flooding.
3. Some river beds now contain a thick layer of plastic getting in the way of fish.
4. Non-biodegradable: not broken­down in the environment by microorganisms.
5. Rubbish is collected and brought to landfill sights causing sight pollution.
6. Plastic is flammable, when burns releases toxic gases.
N.B.
Polymer wastes can be recycled.
• Some are melted down and made into new products.
• Some are melted and cracked to make small molecules that are polymerized into new
plastic but unfortunately some may cannot be recyclable.
• Some are burnet, and heat is used to produce electricity.
What are the advantages and disadvantages of recycling polymers?
• The best long-term solution to disposal would be to work on degradable plastics
• Biodegradable plastics contain some additives as starch that bacteria can feed on.
• Photodegradable plastics contain additives that breakdown in sunlight.
• Biopolymers grow in plants, or made in tanks by bacteria.
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Synthetic Polymers
Synthetic polymers : are polymers are made in factories. Ex. Terylene, lycra,
chewing gum, and plastics as polystyrene and Perspex.
Nylon: [polyamide]
O
O
HO – C –
H
– C – OH + H – N –
H
O
O H
H
–N–H
–C–
–C–N–
–N
+ n H2O
Dicarboxylic acid
+
Diamine
O
polyamide + water
H
—C—N—
(is amide linkage)
­ Nylon is used in making tough, strong fibers which is used in threads, ropes, flying
kits, fishing nets. Parasails and parachutes.
Polyester (PET / polyethylene)
O
HO – C –
O
–C–OH + HO –
– OH
O
O
– C–
–C–O–
–O–
n
+ n H2O
Dicarboxylic acid +
Diol
polyethylene + water
O
—C—O—
(is ester linkage)
­ Polyester is used in making shirts, threads, and polyester­cotton blended fabrics
(more hard wearing than cotton and does not crease so easily).
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Natural Polymers
• Photosynthesis: is a photochemical reaction by which chlorophyll acts as a catalyst
that traps sunlight in presence of H2O and CO2 to produce glucose sugar.
• Plants use glucose as monomers to make starch and cellulose i.e. Carbohydrates.
• Plants use glucose, plus nitrates from the soil to make amino acids.
• Plants use amino acid molecules as monomers to make proteins.
Polymerization
enzymes
Glucose sugar
Starch & Cellulose
Energy store
builds stem
and cell wall
Enzymes
Glucose sugar
Proteins & Fats
• The wood in trees is 50% cellulose
• The polymer in your hair and nails. And in wool and silk and animal horns and
claws, is called keratin.
• The polymer in your skin and bones is called collagen.
Protein:[polyamide / peptide]
Protein is a natural polyamide made of amino acids formed by condensation
reaction and have a wide variety of biological functions.
H
R
N
C
O
C
O H
H
General structure of amino acid where R represents different type of side
chain
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O
H
O
– N –H +HO – C –
HO– C –
Dicarboxylic acid
H
O
–N–H
–C–
H
O
H
–N – C –
–N–
n
Amino acid
+ n H2 O
22 different amino acid are used in building proteins.
• N.B. Protein & Nylon have same amide linkage but different in their monomer.
O
O
HO – C –
– NH2
HO – C –
Amino acid
O
– C – OH + H – N –
dicarboxylic
–N–H
Diamine
Hydrolysis of Polymer
­ Is breaking down of long chain molecules by reaction with water by the effect of
acid, alkali or enzyme.
Hydrolysis
Polymer
Monomers
Reagents for hydrolysis:
• Acidic: HCl(aq) / Δ + H2O
[HCl is a catalyst]
• Alkaline: NaOH(aq) / Δ + H2O
Acidic hydrolysis:
Terylene
Diol + Dicarboxylic acid
Nylon
Diamine + Dicarboxylic acid
Protein
Amino acid
• N.B. Digestion is an acidic hydrolysis.
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