Combined science
chemistry
1
SEPARATION TECHNIQUES
Simple Distillation
Simple Distillation – separation of pure liquid from a solution. Used when we want to obtain the
liquid from the solution.
Process of Distillation:
Solution is heated, and steam (pure vapour) is produced. The steam is cooled in a condenser to form
pure liquid called the distillate. Solute remains in the flask.
Water can be obtained from salt water using this method. The solution is heated in the flask until it
boils. The steam rises into the Liebig condenser, where it condenses back into water. The salt is left
behind in the flask. This sort of technique is used on a much larger scale to obtain pure water for
drinking. This process is carried out in a desalination plant.
Fractional Distillation
Fractional Distillation – separates mixture of miscible (soluble) liquids with differing boiling points.
Use of fractionationg column separates them
Process of Fractional Distillation: E.g. ethanol and water
Mixture of ethanol (b.p 75oC) and water (100oC) is placed in a flask and heated. When the mixture
is heated the vapour produced is mainly ethanol with some steam. Because water has higher boiling
point than alcohol it condenses out from the mixture with ethanol in the fractionating column. The
water condenses and drips back into the flask while the ethanol vapour moves up the column into
the condenser where it condenses into liquid ethanol and is collected in the receiving flask as the
distillate. Temperature will stay constant until all ethanol is distilled. When all the ethanol has
distilled over the temperature begin to rise steadily to 100oC showing that the steam is entering the
condenser. At this point the receiver can be changed and the condensing water can now be
collected.
2
paper chromatography
used to separate two or more dissolved solids in solution.
It makes use of two phases of substances, a moving phase (mobile phase) and non moving phase
(stationery phase).
The stationery phase is the chromatography paper.
The mobile phase is the mixture that must be separated dissolved in a solvent.
A drop of the concentrated solution to be separated is placed on a pencil line (baseline) near the
bottom edge of a strip of chromatography paper as shown below. The paper is then dipped in the
solvent . the level of the solvent must start below the sample. Many solvents are used in
chromatography. Water and organic solvents such as ethanol are common. organic solvents are
useful because they dissolve many substances which are insoluble in water. When an organic
solvent is used the process is carried out in a tank with a lid to stop the solvent evaporating.
The substances separate according to their solubility in the solvent. As the solvent moves up the
paper, the substances are carried with it and begin to separate. The substance that is most soluble
moves fastest up the paper. An insoluble substance would remain at the origin. The run is stopped
just before the solvent front reaches the top of the paper. The solvent front is the furthest point that
the solvent reaches up the chromatography paper. A filter paper can be used as a chromatography
paper.
The distance moved by a particular spot is measured and related to the positon of the solvent front.
The ratio of these distances is called the Rf value.
Rf =
𝑑𝑖𝑠𝑡𝑎𝑛𝑐𝑒 𝑚𝑜𝑣𝑒𝑑 𝑏𝑦 𝑡ℎ𝑒 𝑠𝑢𝑏𝑠𝑡𝑎𝑛𝑐𝑒
𝑑𝑖𝑠𝑡𝑎𝑛𝑐𝑒 𝑚𝑜𝑣𝑒𝑑 𝑏𝑦 𝑡ℎ𝑒 𝑠𝑜𝑙𝑣𝑒𝑛𝑡 𝑓𝑟𝑜𝑛𝑡
3
Chromatography is used to separate both coloured and non-coloured substances. For non-coloured
substances the chromatogram (chromatography paper) is treated with locating agent after the run.
The agent reacts with the samples to produce coloured spots.
Application of paper chromatography
Separation of dyes or extarcts from plants.
MATTER
Structure of an atom
Atoms are made up of sub-atomic particles known as electrons, protons and neutrons. The protons
and neutrons are found in the centre of the atom, which is called the nucleus. The neutrons have
no charge and protons are positively charged. Electrons move around nucleus in an orbit called
electron shells. electrons are negatively charged.
Characteristics of protons, neutrons and protons
Although atoms contain electrically charged particles, the atoms themselves are electrically neutral
(they have no overall electric charge). This is because atoms contain equal numbers of electrons and
protons.
4
Proton number and nucleon number
The number of protons in the nucleus of an atom is called the proton number (or atomic number)
and is given the symbol Z.
Neutrons and protons have a similar mass. Electrons possess very little mass. So the mass of any
atom depends on the number of protons and neutrons in its nucleus.
The total number of protons and neutrons found in the nucleus of an atom is called the nucleon
number (or mass number) and is given the symbol A.
nucleon number = proton number + number of neutrons
(A)
(Z )
For example the element lithium has three protons and four neutrons in its nucleus. It therefore has
a nucleon number of 4+3=7.
The number of neutrons present can be calculated by rearranging the relationship between the
proton number, nucleon number and number of neutrons to give:
number of neutrons = nucleon number − proton number
For example, the number of neutrons in one atom of magnesium with a mass number number of
24 and a proton number of 12 is 24 − 12 = 12
Nuclide notation
The proton number and nucleon number of an element are usually written in the following
shorthand way:
A
where X is the symbol of the element
X
A is the mass number (nucleon number)
Z is the proton number (atomic number)
e.g lithium with a mass number of 7 and a proton number of 3 is written as
3Li
Isotopes
Isotopes are atoms of the same element which have the same proton number but different neutron
numbers. For example carbon has three isotopes i.e carbon 12, carbon 13 and carbon 14 written below
12
6C
13
14
6C
6C
Carbon 12
carbon 13
carbon 14
Carbon has 12-6= 6 nuetrons
Carbon 13 has 13-6= 7nuetrons
Carbon 14 Has 14-6=8 neutrons
The number of protons and electrons remain the same in all the isotopes.
Other examples of isotopes include 16 8O and 18 8O
35
17Cl
and 37 17Cl
Calculate the number of neutrons in each of the isotopes above
5
The arrangement of electrons in an atom
The arrangement of electrons in an atom around the nucleus is known as the electronic
configuration. It is not possible to give the exact position of an electron in an energy level. However,
we can state that electrons can only occupy certain, definite energy levels (shells). Each of the electron
energy shell can hold only a certain number of electrons.



First energy shell holds up to two electrons.
Second energy shell holds up to eight electrons.
Third energy shell holds up to 18 electrons.
The electrons fill the energy levels starting from the energy level nearest to the nucleus, which has
the lowest energy. When this is full (with two electrons) the next electron goes into the second
energy level. When this energy level is full with eight electrons, then the electrons begin to fill the
third and fourth energy levels as stated above.
For example, a 168O atom has a proton number of 8 and therefore has eight electrons. Two of the
eight electrons enter the first energy level, leaving six to occupy the second energy level, as shown
below.
Arrangement of electrons in an oxygen atom.
Other electron arrangement of elements
Hydrogen
Sodium
potassium
H
1
2
9
0
Electronic configuration of the first 20 elements.
Element
Symbol
Proton
number
Number
of
electrons
Electronic
configuration
Hydrogen
H
1
1
1
Helium
He
2
2
2
Lithium
Li
3
3
2,1
6
Beryllium
Be
4
4
2,2
Boron
B
5
5
2,3
Carbon
C
6
6
2,4
Nitrogen
N
7
7
2,5
Oxygen
O
8
8
2,6
Fluorine
F
9
9
2,7
Neon
Ne
10
10
2,8
Sodium
Na
11
11
2,8,1
Magnesium Mg
12
12
2,8,2
Aluminium
Al
13
13
2,8,3
Silicon
Si
14
14
2,8,4
Phosphorus P
15
15
2,8,5
Sulfur
S
16
16
2,8,6
Chlorine
Cl
17
17
2,8,7
Argon
Ar
18
18
2,8,8
Potassium
K
19
19
2,8,8,1
Calcium
Ca
20
20
2,8,8,2
The periodic table
1
2
3
4
5
6
7
8
A period represents elements in same horizontal row e.g from Li to Ne
Elements in a period have the same number of shells.
The period number is equal to the number of shells in an element in that period e.g K (potassium) is
in period 4 therefore has 4 shells, check the structure of potasium drawn earlier above.
A group represents elements in same vertical column e.g from Be to Ra
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The number of electrons in the outer most shell (valency electrons) is equal to the group number e.g
K potassium is in group 1 therefore has 1 electrons in its outer most shell. Check with diagram drawn
above.
Elements in a group have the same number of electrons in the outer most shell (valency electrons).
Thus all elements in group 2 have 2 electrons in their outer most shell.
The elements found between group 2 and group 3 are known as transition elements.
CHEMICAL BONDING
Ionic bonding (electrovalent bonding)
Ionic bonding a usually occur between metals and non-metals.
It gives rise to ionic compounds.
It involves transfer of electrons from the metal atoms to the non-metal atoms during the chemical
reaction. In doing this, the atoms become more stable by getting full outer energy levels.
Metals lose electrons to form positive ions (cations)
Non-metals gain electrons to form negative ions (anions)
For example, consider what happens when sodium and chlorine react together and combine to make
sodium chloride. Sodium has just one electron in its outer energy level (11Na 2,8,1). Chlorine has seven
electrons in its outer energy level (17Cl 2,8,7). When these two elements react, the outer electron of
each sodium atom is transferred to the outer energy level of a chlorine atom as shown in the dot cross
diagram below. In this way both the atoms obtain full outer energy levels and become stable.
Formation magnesium oxide
Magnesium with two electrons in its outer most shell, loses the two electrons. 0xygen with 6 electrons
in its outer shell, gains the two electrons as shown below. In this way both the atoms obtain full outer
energy levels and become stable.
Dot cross diagram for the formation of magnesium oxide
Formation of sodium oxide
An atom of sodium needs to lose one electron to achieve stability but an atom of oxygen needs to
gain two electrons to achieve stability. One in of oxygen therefore needs to bond with two ion of
sodium in order for all ion to achieve stability.
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Dot cross diagram for the formation of sodium oxide
Properties of ionic compounds
 They are usually solids at room temperature
 They have high melting pointsand boiling points. This is due to the strong electrostatic forces
holding the crystal lattice together. A lot of energy is therefore needed to separate the ions.
 They usually cannot conduct electricity when solid, because the ions are not free to move but
they usually conduct electricity when in the molten state or in aqueous solution because the
ions are free to move.
 They mainly dissolve in water. This is because water molecules are able to bond with both the
positive and the negative ions, which breaks up the lattice and keeps the ions apart.
Covalent bonding
it occurs between non metals.
it involves the sharing of a pair electrons by overlapping of orbitals (shells).
For example
Hydrogen (H2)
Each hydrogen atom has one electron, in order to obtain a full outer energy shell and gain the electron
configuration of the noble gas helium, each hydrogen atom must have two electrons. To do this, the
two hydrogen atoms allow their outer energy shells to overlap as shown
Dot cross diagram for hydrogen H2
9
A molecule of hydrogen is formed, with two hydrogen atoms sharing a pair of electrons. This shared
pair of electrons is known as a single covalent bond and represented by a single line as in hydrogen
Dot cross diagram for formation of Chlorine (Cl2)
Water H2O
An oxygen atom has 6 electrons in its outer shell. It therefore needs to share another two in order
to achieve stability. Hydrogen atom has one electron in its outer shell. Therefore two hydrogen
atoms have to bond covalently with one oxygen atom to form water molecule.
Dot cross diagram for formation of water
Properties of covalent compounds
Covalent compounds have the following properties.




As simple molecular substances, they are usually gases, liquids or solids
They have low melting and boiling points. The melting points are low because of the weak
intermolecular forces of attraction which exist between simple molecules.
they do not conduct electricity.This is because they do not contain ions.
they do not dissolve (insoluble) in water but are soluble (dissolve) in organic solvents.
Metallic bonding
Occurs within atoms of metals involving a lattice of positive ions surrounded by a sea of
electrons.
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1. Metals can be bent (ductile) and can be stretched (malleable) because the layers of atoms in
metals slide over each other when force is applied but will not break due to attractive force
between electrons and metal ions.
2. Metals conduct electricity as it has free (delocalised) electrons which carries current.
3. Metals conduct heat as it has free electrons which gains energy when heated and moves faster to
collide with metal atoms, releasing heat in collisions.
4
Metals have high melting and boiling points because the bonds between metals is very strong.
Hence very high heat energy needed to break the bonds.
STOICHIOMETRY
Formula of compounds
the formula shows the ratio of the number of each type of atom that combines to make the
compound.
The elements are represented by their symbols.
Radicals
All radicals may be viewed as ions since they carry either ppositive or negative.
A radical just like an element has a known valency.
Radical
Ionic symbol
Valency
+
Ammonium
NH4
Hydrogen
H+
Nitrate
NO3Chloride
ClHydroxyl
OHIodide
IOxide
Sulphate
Sulphide
Sulphite
Carbonate
Dichromate
O2SO42S2SO32CO32Cr2O72-
Phosphate
PO43-
1
2
3
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Interpreting chemical formula
Consider the following
H means a) hydrogen element symbol
b) one atom hydrogen
H2 implies a) one molecule of hydrogen
b) two atoms of hydrogen element
12
2H2 implies a) two molecules of hydrogen molecules
b) four atoms of hydrogen atoms
similarly it may be deduced that
H2O implies a) one molecule of water
b) two atoms of hydrogen
c) one atom of oxygen
2H2O implies a) two molecules of water
b) four atoms of hydrogen
c) two atoms of oxygen
From the above examples it can be appreciated that the number in front of a formula gives us the
number of molecules of that compound.
The number in front of a formula when multiplied by the number at the bottom right side of the
symbol gives us the total number of atom for that particular.
When a number is after the brackets it affects the number of all atoms within the brackets
e.g Al2(CO3)3
one molecule of aluminium carbonate
two atoms of aluminium
three atoms of carbon
nine atoms of oxygen
Relative atomic mass, Ar
the relative atomic mass of an element is the average mass of its isotope compared with an atom of
carbon 12, 126C.
𝑚𝑎𝑠𝑠 𝑜𝑓 𝑜𝑛𝑒 𝑎𝑡𝑜𝑚 𝑜𝑓 𝑡ℎ𝑒 𝑒𝑙𝑒𝑚𝑒𝑛𝑡
Ar = 𝑚𝑎𝑠𝑠 𝑜𝑓 1⁄
12 𝑜𝑓 𝑎𝑛 𝑎𝑡𝑜𝑚 𝑜𝑓 𝑐𝑎𝑟𝑏𝑜𝑛−12
Relative molecular mass, Mr
It is the sum of the relative atomic masses of all the atoms in the molecule.
Mr (NaOH)
Mr = Ar(Na) + Ar(O) +Ar(H)
=23 + 16 + 1
= 40
Mr (Al2(CO3)3) = Ar(Al) x 2 + Ar(C) x 3 + Ar(O) x 9
= (27 x 2) + (12 x 3) + (16 x 9)
= 234
The mole
The mole is used to replace the terms atomic mass unit and molecular mass unit.
One mole is the
1. Amount of substance which contains an Avogadro number 6.02 x 1023
1 mole of oxygen = 6.02 x 1023
1 mole of potassium = 6.02 x 1023
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2. Relative atomic mass, Ar of an element in grams (g)
1 mole of lithium = Ar(g)
= 7g
1 mole of sulphur = Ar(g)
= 32g
3. Relative molecular mass, Mr of a compound or molecule in grams (g)
1 mole of water = Mr(H2O)
=Ar(H) x 2 + Ar(O)
= (1 x 2) + 16
=18g
1 mole of sodium sulphide = Mr(Na2S)
= Ar(Na) x 2 + Ar(S)
= (23 x 2) + 32
=78g
questions
Calculate the mass of 5 moles of calcium chloride, CaCl2
1.
𝑀
N=𝑀𝑟
Where n is the number of moles
M is the mass in grammes (g)
Mr is the relative molecular mass
𝑀
5= 111
M = 5x 111
=555g
2. How many moles are in 72g of water
𝑀
N=𝑀𝑟
72
N=18
N= 4moles
Chemical equations
Characteristics
1. A chemical equation is a short and convenient way of describing a chemical reaction.
2. The total number of all atoms involved do not change nor disappear, they are just
rearranged
3. The number of different atoms are the same on both sides of the equations. This is called a
balanced.
4. States of the reactants should always be added to a chemical equation. The state symbols
are (s) for solid
(l) for liquid
(g) for gas
(aq) for aqueous
14
writing a chemical equation
the following steps can be used
1. write the word equation
magnesium + oxygen → magnesium oxide
2. write the correct corresponding formulas
Mg + O2 → MgO
of the chemicals involved
3. Check if equation is balanced. If not balanced,
balance the equation atom by atom.
2Mg + O2 → 2MgO
Remember never to balance by adding number
To the right of the symbol.
4. Add state symbols to the right of the formula
2Mg(S) + O2(g) → 2MgO(s)
Examples
Zinc + chlorine → zinc chloride
Zn + Cl2 → ZnCl2
Calcium + water → calcium hydroxide + hydrogen
Ca + H2O → Ca(OH)2 + H2
Ca + 2H2O → Ca(OH)2 + H2
The Avogadro number
The number of particles contained in one mole of any substance is called the Avogadro number. This
number is 6.02 x 1023/mol and is given the symbol L.
1mole of carbon contains 6.02 x 1023 carbon atoms
1 mole of sodium contains 6.02 x 1023 sodium atoms
1 mole of carbon dioxide contains 6.02 x 1023 carbondioxide molecules
e.g how many grams of magnesium contain the Avogadro number of particles 6.02 x 1023
1 mole of magnesium is 24g and
1 mole of magnesium is 6.02 x 1023
Therefore 24g of magnesium contain the avogadro particles.
e.g How many particles are in 0.2 moles of carbon
1 mole of carbon has 6.02 x 1023 particles
0.2moles has less
0.2
1
x 6.02 x 1023
1.2 x1023particles
concentration
concentration is a measure of how much solute is dissolved in a solvent.
The S.I unit of concentration are g/dm3 or mole/dm3
Note 1dm3 = 1000cm3 = 1000ml =1L
The following relationship can be used to calculate concentration
n = VC
15
where n is the number of moles
V is the volume in dm3
C is the concentration in mol/dm3
Also note that the molarity of a solution is the number of moles of solute present in 1dm3 of a
solution
Therefore mol/dm3 is often shortened to M, e.g 0.1 mol/dm3 is written as 0.1M (M is molar).
1. Find the volume of 0.1 mol/dm3 of potassium hydroxide solution that contains 0.2
mol/dm3of potassium hydroxide
n = vc
𝑐
v=𝑛
0.2
v = 0.1
v = 2dm3
2. if 8g of sodium hydroxide are dissolved to form 2500cm3 of sodium hydroxide, calculate its
concentration.
n = vc
first find the number of moles, n =
=
𝑀
𝑀𝑟
8
40
= 0.2moles
𝑛
c=𝑣
c=
0.2
2.5
= 0.08mol/dm3
Empirical formula
Empirical formula shows the simplest ratio of the atoms present.
To find the empirical formula
i. divide each percentage by the atomic mass
ii. divide the ratios obtained in (i) by the smallest numbers.
1. In another experiment an unknown organic compound was found to contain 0.12 g of
carbon and 0.02 g of hydrogen. Calculate the empirical formula of the compound. (Ar: H = 1;
C = 12)
Masses (g)
Number of moles
C
0.12
H
0.02
0.12
0.02
12
1
16
Ratio of moles
Empirical formula
= 0.01
= 0.02
0.01
0.02
0.01
0.01
1
CH2
2
2. Calculate the empirical formula of an organic compound containing 92.3% carbon and 7.7%
hydrogen by mass. The Mr of the organic compound is 78. What is its molecular formula? (Ar: H = 1;
C = 12)
C
92.3
% by mass
H
7.7
92.3
Number of moles
7.7
12
1
= 7.7
Ratio of moles
Empirical formula
= 7.7
7.7
7.7
7.7
7.7
1
CH
1
Molecular formula
Molecular formula shows the actual numbers of atoms of each element present in the molecule.
1. Calculate the empirical formula of an organic compound containing 80% carbon and 20% hydrogen
by mass. The Mr of the organic compound is 30. What is its molecular formula? (Ar: H = 1; C = 12)
first find the empirical formula of the compound
C
H
% by mass
80
20
Number of moles
80
20
12
1
= 6.67
Ratio of moles
Empirical formula
= 20
6.67
20
6.67
6.67
1
3
CH3
Mr( CH3)n =30
[Ar(C) x n] + [Ar(H) x 3 x n] = 30
(12 x n) + (1 x 3 x n) = 30
12n + 3n = 30
15n = 30
30
n = 15
n=2
(CH3)2 = C2H6
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ACIDS AND BASES
Properties of acids and bases
Property
Taste
Acid
Sour.
base
Bitter.
Texture
Sticky.
Soapy.
Litmus paper
Blue litmus turns red.
Red litmus turns blue.
Universal indicator
Turns yellow, orange, red
depending on strength of
acid.
Turns blue, purple depending
on strength of base.
reactions
Reacts with bases, metals,
carbonates to form salts.
Reacts with acids to form salts.
pH
pH is less than 7
pH is above 7
Examples
Corrosive
Hydrochloric acid, sulphuric
acid, vinegar, citric acid
corrosive
Ammonia, potassium
hydroxide, sodium hydroxide
pH scale
it is a measure of how acidic or alkaline a substance is.
The pH scale ranges from 0 to 14.
Acids have a pH below 7.
Bases have a pH above 7.
A pH of is neutral, it is neither acidic or basic
Indicators
an indicator is asubstance that changes colour according to the acidity or alkalinity of the solution it
is in. indicators are used to measure the pH of a solution. The table shows a few indicators and their
colour changes
Indicator
Colour in acid
Colour in base
Litmus
Red
Blue
Methyl orange
Red
Yellow
Phenolphthalein
Colourless
Red
Universal indicator
Red in strong acid
Purple in strong base
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Strong and weak acids
An acid is a substance that dissociates to produce H+ ions.
Strong acids are substances that completely dissociate into ions. Eg sulphuric acid (H2SO4) and
hydrochloric acid (HCl). The term strong refers to extent of dissociation and not the concentration.
Adding water to an acid only dilutes it but does not make it weaker acid.
Weak acids are substances that partial dissociate eg ethanoic acid, vinegar.
Strong and weak alkali
A base is a substance that dissociates to produce OH- ions.
Bases are oxides or hydroxides of metals
Alkalis are bases which are soluble in water
Strong alkalis are substance that completely dissociate in water. Eg sodium hydroxide (NaOH) and
Potassium hydroxide (KOH).
Weak alkalis are substances that partial dissociates in water. Eg ammonia (NH3)
Reaction of acids with metals
Acids react with metals to produce a salt and hydrogen gas. The gas is tested with a burning splint
which shows hydrogen burns with a ‘pop’ sound. The general equation is
The salt produced has two words and will depend on the metal and the acid used. The first word
is the name of the metal used while the last word is depended upon the acid used in the reaction.
If sulphuric acid (H2SO4) is used it becomes sulphate
If hydrochloric acid (HCl) is used it becomes chloride
If nitric acid (HNO3) is used it becomes nitrate
If ethanoic acid (CH3COOH) is used it becomes ethanoate
e.g
sodium + hydrochloric acid → sodium chloride + hydrogen
2Na(s) + 2HCl(aq) → 2NaCl(aq) + H2(g)
Reaction of acids with carbonates
- Acids react with carbonates to produce a salt, carbon dioxide gas and water. Carbon dioxide
produced is tested by bubbling into limewater which forms a white precipitate. The general
equation is
Acid + metal carbonate →salt + carbon dioxide + water
e.g sulphuric acid + magnesium carbonate → magnesium sulphate + carbon dioxide + water
H2SO4(aq) + MgCO3(s) → MgSO4(aq) + CO2(g) + H2O (l)
- Reaction of Acids with bases (neutralisation)
Acid react with bases or alkalis to form salt and water. The general equation is
Acid + base →salt +water
e.g sulphuric acid + copper hydroxide → copper sulphate + water
H2SO4(aq) + Cu(OH)2(s) → CuSO4(aq) + 2H2O (l)
Note: neutralisation reaction is the reaction between an acid and base to form salt and water.
19
Acid-Base titrations
Titration is the slow addition of a solution of known concentration to a given volume of another
solution of unknown concentration to a point of neutralisation. The solution of known concentration
is called the titrant and the solution of unknown concentration is called the analyte.
Apparatus used in titration
Burette, clamp and stand, conical flask, volumetric pipette, dropper or small pipette and indicator.
Performing a titration
1.
2.
3.
4.
5.
6.
7.
8.
9.
Write a balanced equation for the reaction.
Measure required volume of analyte using a pipette and run it into a conical flask.
Add a few drops of the indicator into the conical flask.
Fill the titrant into the burette. Zero the burette by opening the tap letting the titrant into a
beaker until the meniscus is level with the maximum mark on the burette. N.B- this is
necessary to get rid of any air bubbles in the titrant.
Place the conical flask below the burette so that the tip of the burette is well within the
conical flask.
Slowly open the tap of the burette just enough to let a few drops into conical flask.
Continue to let a few drops of titrant a little at a time, reducing volume added as you
progress. Also swirl contents of flask after each addition.
As soon as the indicator starts to change colour, close the tap of burette immediately and
swirl the mixture making sure the colour change is final. This shows that all the analyte has
reacted. Addition of more results in wrong results.
Record volume of titrant remaining in burette.
20
10. Calculate volume used by subtracting end point volume from starting volume.
Acid-base titrations
This is a neutralisation titration, whose aim is to produce a salt which is separated and dried.
Experiment
Aim: to use an acid base titration to produce a salt.
Materials: Burette, clamp and stand, conical flask, 25ml pipette, dropper or small pipette,
phenolphthalein indicator, dilute sodium hydroxide, evaporating dish, dilute HCl, Bunsen burner.
Procedure
1. Write a balanced equation for the reaction.
2. Measure 25ml of dilute sodium hydroxide (analyte) using a pipette and run it into a
conical flask.
3. Add a few drops of the phenolphthalein indicator into the conical flask.
4. Fill the dilute HCl (titrant) into the burette. Zero the burette by opening the tap letting the titrant
into a beaker until the meniscus is level with the maximum mark on the burette. N.B- this is
necessary to get rid of any air bubbles in the titrant.
5. Place the conical flask below the burette so that the tip of the burette is well within the conical
flask.
6. Slowly open the tap of the burette just enough to let a drop into conical flask.
21
7. Continue to let a drop of titrant a little at a time, reducing volume added as you progress. Also
swirl contents of flask after each addition.
8. As soon as the indicator starts to change colour, close the tap of burette immediately and swirl
the mixture making sure the colour change is final. This shows that all the dilute sodium hydroxide
(analyte) has reacted. Addition of more; results in wrong results.
9. Record volume of titrant remaining in burette.
10. Calculate volume used by subtracting end point volume from starting volume.
11. Pour the neutral solution into evaporating dish and heat so that excess water evaporates.
12. Leave the saturated solution to cool and crystallise.
INDUSTRIAL PROCESSES
Liquefaction and Fractional distillation of air
• The air is passed through fine filters to remove dust.
• The air is cooled to about −80 °C. water vapour and carbon dioxide solidify and are removed. If
these are not removed, then serious blockages of pipes can result.
• Next, the cold air is compressed to about 100 atm of pressure. This warms up the air, so it is
passed into a heat exchanger to cool it down again.
• The cold, compressed air is allowed to expand rapidly, which cools it still further.
• The process of compression followed by expansion is repeated until the air reaches a temperature
below −200 °C. At this temperature the majority of the air liquefies
• The liquid air is passed into a fractionating column and it is fractionally distilled. The gases can be
separated because they have different boiling points. Nitrogen boils of at -196oC and oxygen at 1830C.
• The gases are then stored separately in large tanks and cylinders.
Uses of the gases
Gas
Oxygen
Nitrogen
Carbondioxide
Use
Welding
Steel making
Medical purposes
Basic oxygen furnace
Make ammonia and nitric acid
Refrigerant
Food preservative
-Used in fire extinguishers
-As dry ice or coolant.
-Used in fizzy drinks as a preservative.
Electrolysis
Electrolysis is the decomposition of compound using electricity.
Electrolysis is an electrochemical reaction in which electricalenergy is converted to chemical energy.
Electrolysis takes place in an electrolytic cell as shown below;
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Electrolyte is an ionic compound which conducts electric current in molten or aqueous solution, being
decomposed in the process.
Electrode is a rod or plate where electricity enters or leaves electrolyte during electrolysis. Reactions
occur at electrodes. Electrodes usually made of unreactive metals such as platinum or of the nonmetal
carbon (graphite). These are said to be inert electrodes because they do not react with the products
of electrolysis. The names given to the two electrodes are cathode, the negative electrode (connected
to the negative terminal of the battery) which attracts cations (positively charged ions), and anode,
the positive electrode (connected to the positive terminal of the battery) which attracts anions
(negatively charged ions).
Principles of electrolysis
If electricity is passed through an electrolyte, the electrolyte breaks down to form ions.
The positively charged ions move to the negative electrode (cathode). They receive electrons and are
reduced. Thus reduction always takes place at the cathode.
The negatively charged ions move to the positive electrode (anode). They lose electrons and are
oxidised. Thus oxidation always take place at the anode.
Electrolysis of molten lead (II) bromide
To make molten lead(II) bromide, PbBr2, we strongly heat the solid until it melts. To electrolyse it, pass
current through the molten PbBr2.
Ions Present
Pb2+ and BrThe Pb2+ ions attracted (move) to the cathode while the Br- ions are attracted (move) to the anode.
23
Reaction at Anode
Br- (bromide ions) loses electrons at anode to become Br (bromine) atoms. Br atoms created form
bond together to make Br2 gas.
2Br- → Br2 + 2e-
Reaction at Cathode
Pb2+ (lead II ions) gains electrons at cathode to become Pb (lead) atoms.
Pb2+ + 2e- → Pb
Overall Equation
PbBr2 → Pb + Br2
Products of electrolysis of lead bromide
The products are lead metal and bromine gas
Experiment; electrolysis of lead bromide
set up apparatus as shown;
results and explanation
When the lead bromide is solid the bulb does not light, because the solid compound does not allow
electricity to pass through it.
However, when the compound is heated until it is molten, the bulb does light because the molten
lead(ii) bromide now has free ions therefore behaves as an electrolyte.
When this happens bubbles are seen at the anode and lead metal is produced at the cathode
underneath.
When the heat is removed everything stops. The bulb goes off and the bubbles stop being produced.
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Electrolysis of water
Water is a poor conductor of electricity. To make it conduct, a little dilute acid is added (dilute
sulphuric acid).
Ions Present
H+, OH- and SO42H+ ions moves to the cathode.
OH- and SO42- moves to the anode. At the anode the OH- ions are discharged
Reaction at Anode
OH- loses electrons at anode to become O2 and H2O.
4OH -(aq) → O2(g)+ 2H2O(l) +4e-
Reaction at Cathode
H+ gains electrons at cathode to become H atoms becoming hydrogen gas.
2H+(aq) + 2e- → H2(g)
Overall Equation
2H2O(l)→ 2H2(g) + O2(g)
Since only water is electrolysed, the sulphuric acid now only becomes concentrated.
Products of electrolysis of water
The products are hydrogen and oxygen.
The amount of hydrogen produced during electrolysis of water is twice the amount the amount of
oxygen. The ratio of hydrogen to oxygen is 2:1.
Uses of hydrogen
Welding
Manufacture of margarine
Manufacture of ammonia
Manufacture of fertilisers
Electrolysis of CuSO4 Using Inert Electrodes(e.g. carbon)
Ions Present
Cu2+, H+, OH- and SO4225
Reaction at Anode
OH- loses electrons at anode to become O2 and H2O.
4OH -(aq)→ O2(g)+ 2H2O(l) +4e-
Reaction at Cathode
Cu2+ gains electrons at cathode to become Cu atoms becoming liquid copper.
Hydrogen ions are not discharged because copper is easier to discharge.
Cu2+(aq) + 2e- →Cu(s)
Since copper ions in solution are used up, the blue colour fades. Hydrogen and sulphate ions left forms
sulphuric acid.
Electrolysis of CuSO4 Using Active Electrodes(e.g. copper)
Ions Present
Cu2+, H+, OH- and SO42-
Reaction at Anode
The copper anode lose electrons to form Cu2+. So, the electrode size decreases.
Cu(s)→ Cu2+(aq) + 2e-
Reaction at Cathode
Cu2+ produced from anode gains electrons at cathode to become Cu atoms becoming copper. Hence,
the copper is deposited here and the electrode grows.
Cu2+(aq) + 2e- →Cu(s)
Overall Change
There is no change in solution contents as for every lost of Cu2+ ions at cathode is replaced by Cu2+
ions released by dissolving anode. Only the cathode increases size by gaining copper and anode
decreases size by losing copper. We can use this method to create pure copper on cathode by using
pure copper on cathode and impure copper on anode. Impurities of anode falls under it.
Electroplating
Electroplating is coating an object with thin layer of metal by electrolysis. This makes the object
protected and more attractive.
26
Object to be plated is made to be cathode and the plating metal is made as anode.
The electrolyte MUST contain plating metal cation.
Plating Iron object with Nickel
Reaction at Anode
Ni2+ discharged from anode into solution. So, the electrode size decreases.
Ni(s) → Ni2+(aq) + 2e-
Reaction at Cathode
Ni2+ produced from anode gains electrons at cathode to become Ni atoms becoming nickel. Hence,
the nickel is deposited here and the electrode grows.
Ni2+(aq) + 2e- → Ni(s)
Copper plating an iron nail
The anode is copper electrode
The cathode is the iron nail
The electrolyte is copper sulphate
Reaction at Anode
The copper anode lose electrons to form Cu2+. So, the electrode size decreases.
Cu(s) → Cu2+(aq) + 2e-
Reaction at Cathode
Cu2+ produced from anode gains electrons at cathode to become Cu atoms becoming copper. Hence,
the copper is deposited here and the electrode grows.
Cu2+(aq) + 2e- →Cu(s)
Reasons for electroplating
Decorative purposes
Prevent corrosion
OXIDATION AND REDUCTION
Reduction-Oxidation in terms of oxygen
Oxidation is the gain of oxygen by a substance
Reduction is the loss of oxygen by a substance
e.g Pb(s) + Ag2O (aq) → PbO(aq) + 2Ag (aq)
This reaction involves both oxidation and reduction therefore is called a redox reaction. A redox
reaction is whereby oxidation and reduction occur simultaneously in a reaction.
27
From the reaction;
Pb is oxidized as it gains oxygen from Ag2O to form PbO.
Ag2O is oxidizing agent.
Ag2O is reduced as it loses oxygen to Pb to form Ag.
Pb is reducing agent.
Oxidizing agent is a substance which causes oxidation of another substance
Reducing agent is a substance which causes reduction of another substance
Reduction-Oxidation in terms of hydrogen
Oxidation is the loss of hydrogen by a substance
Reduction is the gain of hydrogen by a substance
e. g
H2S(g) + Cl2(g) → 2HCl(g) + S(g)
H2S is oxidized as it loses hydrogen to Cl2 to form S.
Cl2 is oxidizing agent.
Cl2 is reduced as it gains hydrogen from H2S to form HCl.
H2S is reducing agent.
Reduction-Oxidation in terms of electron transfer
Oxidation is the loss of electrons by an atom
Reduction is the gain of electrons by an atom
e.g
2Na(s) + Cl2(g) → 2NaCl(s)
inorder to identify the oxidised and reduced substance, it is important to rewrite the equation showing
ions present in compounds involved in the reaction. e.g the above equation is rewritten as
2Na(s) + Cl2(g) → 2Na+Cl-(s)
It becomes clear which atom has gained or lost electrons.
When an atom becomes positive, more positive or less negative, it has lost electrons, thus has been
oxidised.
When an atom becomes negative, more negative or less positive, it has gained electrons, thus has
been reduced.
From the above reaction
Na becomes positive, Na+. Therefore Na is oxidized as it loses electron to Cl2 to form Na+ ions.
Cl2 is oxidizing agent.
Cl2 becomes negative, Cl-. Therefore Cl2 is reduced as it gains electron from Na to form Cl- ions.
Na is reducing agent
Oxidation State in Reduction-Oxidation reaction
Oxidation State is the charge an atom would have if it existed as an ion.
To work out oxidation state, the rules are:
1. all elements in their natural state have oxidation state zero, e.g. Cu, Fe, N2
2. Oxidation of an ion is the charge of the ion, e.g. Na+ = +1, Cu2+=+2, O2- = -2
3. The oxidation state of some elements in their compounds is fixed, e.g.
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Group I Elements = +1
Group II Elements = +2
Hydrogen in most compounds = +1
4. The oxidation states of elements in a compound add up to 0
NaCl: (+1) + (-1) = 0
K2O: (+1) x 2 + (-2) = 0
.
Al2O3: (+3) x 2 + (+2) x 3 = 0
5. Some compounds with possible variable oxidation states have roman numeral as a guide about
their oxidation state, e.g
- Iron(II) chloride has formula FeCl2 and iron oxidation state +2
Potassium(VI) dichromate has formula K2Cr2O7 and potassium oxidation state +6
- Manganese(IV) oxide has formula MnO2 and manganese oxidation state +4
6. Sum of oxidation states of elements in an ion is equal to charge on the ion, e.g.
OH-: (-2) + (+1) = -1
Examples:
Work out the oxidation states of the underlined elements in these compounds:
(a) CO2
(oxidation state of C) + (-2) x 2 = 0
(oxidation state of C) + (-4) = 0
Oxidation state of C = +4
(b) KMnO4
(+1) + (oxidation state of Mn) + (-2) x 4 = 0
(oxidation state of Mn) + (+1) + (-8) = 0
(oxidation state of Mn) + (-7) =0
Oxidation state of Mn = +7
(c) Fe(NO3)2
(oxidation state of Fe) + (-1) x 2 = 0
(oxidation state of Fe) + (-2) = 0
Oxidation state of Fe = +2
Oxidation and reduction in terms of oxidation number
Oxidation is the increase of oxidation state by a substance
Reduction is the decrease of oxidation state by a substance
With oxidation numbers it is also important to rewrite the equation showing ions present in
compounds involved in the reaction.
e.g Cu(s) + HCl(aq) → CuCl2(aq) + H2(g)
Cu(s) + H+Cl-(aq) → Cu2+Cl-(aq) + H2(g)
Cu is oxidized as it gains oxidation state from 0 to +2.
Cu is reducing agent
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H+ions in HCl reduced as it loses oxidation state from +1 to 0.
H+ions are oxidising agent
Nothing has happened to Cl since there was noo change in oxidation state
Another example
Cl2(aq) + 2KI(aq) → 2KCl(aq) + I2(aq)
Cl2(aq) + 2K+ I- (aq) → 2K+ Cl- (aq) + I2(aq)
I-ions in KI is oxidized as it gains oxidation state from -1 to 0.
I-ions is reducing agent
Cl2 is reduced as it loses oxidation state from 0 to -1.
Cl2 is oxidizing agent
Extraction of iron
The main ores of iron are haematite and magnetite
Iron is extracted in a blast furnace
Raw materials
Iron ore (Fe2O3)
Coke (C)
Limestone (CaCO3)
The process
The raw materials are mixed together to form what is called charge. The charge is fed into the top of
the blast furnace
1. A blast of hot air is sent in near the bottom of the furnace through holes (tuyères) which makes
the ‘charge’ glow, as the coke burns in the preheated air.
C(s)+ O2(g) → CO2(g)
2. The carbon dioxide gas produced reacts with more hot coke higher up in the furnace, producing
carbon monoxide in an endothermic reaction.
carbon dioxide + coke →carbon monoxide
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CO2(g) + C(s) →2CO(g)
3.
Carbon monoxide is a reducing agent. it rises up the furnace and reduces the iron
(III)oxide ore. This takes place at a temperature of around 700 °C:
iron(III) oxide + carbon monoxide → iron + carbon dioxide
Fe2O3(s) + 3CO(g) → 2Fe(l) + 3CO2(g)
The molten iron trickles to the bottom.
4. Because of high temperatures in the furnace the limestone (CaCO3) decomposes.
Calcium carbonate → calcium oxide + carbon dioxide
CaCO3(s) → CaO(s) + CO2(g)
5.
Calcium oxide reacts with impurities such as silicon(iv) oxide in the iron, to form a slag which is
mainly calcium silicate.
calcium oxide + silicon(iv) oxide →calcium silicate
CaO(s) +SiO2(s) → CaSiO3(l)
The slag trickles to the bottom of the furnace, but because it is less dense than the molten iron, it
floats on top of it.
The molten iron, as well as the molten slag, may be tapped off (run off) at regular intervals.
The waste gases, mainly nitrogen and oxides of carbon, escape from the top of the furnace. They are
used in a heat exchange process to heat incoming air and so help to reduce the energy costs of the
process. Slag is the other waste material. It is used by builders and road makers for foundations.
The iron obtained by this process is known as ‘pig’ or cast iron and contains about 4% carbon (as well
as some other impurities.
Production of steel
The ‘pig iron’ obtained from the blast furnace contains between 3% and 5% of carbon and other
impurities, such as sulfur, silicon and phosphorus. These impurities make the iron hard and brittle. In
order to improve the quality of the metal, most of the impurities must be removed and in doing this,
steel is produced.
The impurities are removed in the basic oxygen process, which is the most important of the
steelmaking processes. In this process:
Molten pig iron from the blast furnace is poured into the basic oxygen furnace
A water-cooled ‘lance’ is introduced into the furnace and oxygen is blown onto the surface of the
molten metal.
Carbon is oxidised to carbon monoxide and carbon dioxide, while sulfur is oxidised to sulfur dioxide.
These escape as gases.
Silicon and phosphorus are oxidised to silicon(iv) oxide and phosphorus pentoxide, which are solid
oxides.
Some calcium oxide (lime) is added to remove these solid oxides as slag. The slag may be skimmed or
poured off the surface.
Samples are continuously taken and checked for carbon content. When the required amount of
carbon has been reached, the blast of oxygen is turned off.
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There are various types of steel that differ only in their carbon content. The differing amounts of
carbon present confer different properties on the steel and they are used for different purposes.
Alloys of iron
Steel
Typical composition
Properties
Uses
Mild steel
99.7% iron, 0.3% carbon
Easily worked
Lost most of brittleness
Low cost
Car bodies, large structures,
machinery, railway tracks,
blocks and bolts
Hard steel
99% iron, 1% carbon
Tough and brittle
Cutting tools, chisels, razor
blades
Manganese
steel
87% iron, 13% manganese
Tough, springy
Drill bits, springs
Stainless
steel
70% iron, 20% chromium,
10% nickel
Tough, does not corrode
Cutlery, kitchen sinks, surgical
instruments, chemical plant
reaction vessels, car parts
Tungsten
steel
82-99% iron,
tungsten
Tough, hard, even at high
temperatures
high-speed
cutting
and
sharpening tools, machinery,
crushing machinery.
Wrought
iron
99% iron, 1% slag
Soft but it does not melt can
be worked into shapes using
tools.
Ornate railings and gates,
garden furniture
Cast iron
96-98% iron, 2-4%carbon
Hard and brittle with a low
melting point;
easy to melt and mould
Cooking utensils
decoration
Medium
steel
99.5% iron, 0.5% carbon
Tougher than mild steel
Car springs, chains
Havel (high
carbon
steel)
99% iron, 1% carbon
Very hard but brittle
Blades for cutting tools, chisel
and razor blades
Nickel steel
95%iron, 5% nickel
1-18W%
hard and tough
Turbine blades, axles, storage
canisters for liquefied gas
NOTE: AN ALLOY IS A SUBSTANCE MADE UP OF TWO OR MORE METALS CHEMICALLY COMBINED
The Manufacture of Ammonia: The Haber Process
Ammonia is produced by the Haber process. Ammonia is a poisonous colourless gas with a pungent
smell.
Raw materials
Nitrogen; obtained from fractional distillation of liquid air.
Hydrogen; obtained from electrolysis of water.
The Process
Nitrogen and hydrogen are mixed together in ratio 1:3 and passed over iron catalyst.
The reaction is
32
Since the reaction is reversible so H2 and N2, reproduced from decomposition of produced NH3, are
passed over the catalyst again to produce ammonia.
Conditions for Manufacturing Ammonia
Temperature of 450oC
Pressure of 200atm
Iron catalyst
Explaining the conditions
Graph shows that to have high yield of ammonia we should have:
1. Higher pressure
2. Lower temperature
But in practice, we use lower pressure of 200 atm because using high pressure involves safety risk and
higher cost and higher temperature of 450oC because using low temperature is too slow to reach
equilibrium
Uses of ammonia
For producing: fertilisers, nitric acid, nylon, dyes, cleaners and dry cell.
Manufacture of sulphuric acid; the contact process
Sulphuric acid is manufactured by the contact process
Raw materials
Sulphur /sulphide ores
Air
Water
33
The contact process
1. Production of sulphur dioxide SO2
Sulphur or sulphide ore is burnt in air to produce sulphur dioxide, SO2
S(g) + O2(g) → SO2(g)
2. Sulphur dioxide is further reacted with oxygen at a temperature between 4500C and
5000C, pressure of 1atm and in the presence of vanadium (V) oxide (V2O5) catalyst to
produce sulphur trioxide.
3. Sulphur trioxide is cooled and reacted in concentrated H2SO4 to produce oleum, H2S2O7.
Sulphur trioxide, SO3 is not reacted with water as mist forms which is difficult to condense.
SO3 + H2SO4 → H2S2O7
4.
Oleum is diluted with water to produce sulphuric acid.
H2S2O7 + H2O → 2H2SO4
Optimum condition for the contact process
temperature between 4500C and 5000C,
pressure of 1atm
vanadium (V) oxide, (V2O5) catalyst
explaining the conditions
Since the reaction
is a reversible and exorthermic reaction, a
low temperature and high pressure is needed to enhance the forward reaction.
Although lower temperature is required for high yield, 4500C is instead used in the process as the
reaction will be too slow with low temperatures and high temperatures will decompose the
sulphur trioxide. This is the reason why 4500C is called a compromise temperature.
A catalyst vanadium(V) oxide, V2O5, is used to increase the yield of sulphur trioxide. Platinum can
catalyse more efficiently but it is too expensive.
Uses of Sulphuric Acid
- Making of fertilizers such as superphosphate and ammonium sulphate
- Making detergents
- Cleaning surfaces of iron and steel surface before galvanization or electroplating
- To manufacture plastics and fibres
- As electrolyte in car batteries
- In refining of petroleum
- In production of dyes, drugs, explosives, paints, etc
Manufacture of nitric acid
raw materials
ammonia from haber process
air
water
34
the process
ammonia is mixed with air and is passed over a red hot platinum-rhodium catalyst.
4NH3 + 5O2 → 4NO + 6H2O
On cooling the nitrogen monoxide (NO) reacts with more oxygen from the air and is further oxidised
to nitrogen dioxide( NO2).
2NO + O2 → 2NO2
This is finally absorbed in water in the presence of oxygen
4NO2 +2 H2O + O2 → 4HNO3
Manufacture of ammonium nitrate fertiliser
Raw materials
ammonia gas
nitric acid.
The process
Ammonia is reacted with nitric acid
ammonia + nitric acid → ammonium nitrate
NH3(g) + HNO3(aq) → NH4NO3(aq)
ORGANIC CHEMISTRY
Hydrocarbons
Hydrocarbons are organic compounds containing hydrogen and carbon atoms only e.g alkanes and
alkenes.
Homologous series is a series of compounds that have similar properties and the same general
formula but hose structure only differ by the number of CH 2 units in the main carbon chain. For
example alkanes are a homologous series and alkenes too as well as alcohols.
Identifying homologous series
Organic compounds have a functional group which enables them to be identified to a particular
homologous series. The following table gives the functional groups.
Homologous series
Functional group
Alkanes
C—C
alkenes
C=C
alcohols
—OH
Naming organic compounds
Organic compounds names have a prefix and a suffix.
The prefix is obtained from the number of carbon atoms in the longest chain.
Number of carbons Prefix
1
Meth2
Eth3
Prop4
But5
Pent-
35
The suffix is obtained from the homologous series or functional group
Homologous series
Functional group
Suffix (name ending)
Alkanes
C—C
-ane
alkenes
C=C
-ene
alcohols
—OH
-anol
(-ol)
How to name organic compounds
1. Identify the prefix by counting the number of carbon atoms in the chain.
2. Identify the suffix from the functional group
e.g name the compound
1. It has one carbon atom, meth2. It is an alkane, -ane
Therefore the name is methane
1. It has two carbon atoms, eth2. It is an alkane, -ane
Therefore it is ethane
1. It has two carbon atoms, eth2. It is an alkene, -ene
therefore it is ethene
Drawing organic compounds
The prefix gives you the number of carbon atoms in the molecule.
The suffix gives the functional group.
e.g draw the structure of propane
1. the prefix prop- means it has 3 carbon atoms
2. the suffix –ane means it is an alkane (C—C)
Therefore the structure is
The structure of propene
1. The prefix pro- means it has three carbon atoms
2. The suffix –ene means it is an alkene
36
therefore the structure is
Production of biogas
Biogas is a gas evolved when organic material decomposes. It is a mixture of gases including
methane, carbon dioxide and hydrogen sulphide.
Biogas is obtained from fermenting manure.
the manure is put into a sealed pit. Bacteria digest the manure and the gas given off is
collecte and piped to burners. The sludge in the pit has to be agitated by rotating the
floating gasometer tank on the top several times a day.eventually the sludge build up has to
be removed and the process started again. The sludge can be used as a fertilizer for crops.
The biogas plant/bio-digester
Factors affecting the production of biogas
 Temperature- optimum temperatures are between 350C and 550C
 P.H- the solution should not be too acidic or too alkaline
 Role of yeast
 Type of waste
 Time
Uses of Biogas
-For cooking
-For lighting
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Ethanol
Ethanol is an organic compound which belongs to the homologous series, alcohol.
The structure of ethanol is
Ethanol is a liquid fuel that burns with a pale blue flame and leaves no residue. It is made by
fermenting plants that have high sugar content / maize. In Zimbabwe, ethanol is produced
at the ethanol plant at Triangle.
Ethanol from sugar cane
The cane is crushed and the juice from it. The fibre that remains is called bagasse and is
used as fuel for the boilers that generate steam. The juice is filtered and evaporated to
produce sugar crystals. These are separated from the molasses.
The molasses from the sugar mill is diluted with water and mixed with a yeast culture. The
pH is lowered by adding sulphuric acid. The yeast will survive at this pH but bacteria cannot.
Air is bubbled through the tank and the yeast multiplies.
When the yeast population is at the required level, the culture is removed to the main
fermentation vats. Air is not admitted at this stage so that the fermentation will be
anaerobic. Under these conditions yeast cannot multiply. The pH is adjusted and the
mixture is allowed to brew for 40hours.
The enzymes in yeast cause the breakdown of the sugars into smaller glucose molecules
which is then converted into ethanol by the action of zymase in yeast. Carbon dioxide is the
other product.
Once the alcohol content reaches about 15%, the yeast becomes inactive. Fractional
distillation is then used to concetrate the alcohol.
Fractional Distillation of ethanol
38
Note: one percent Benzol is added to the pure ethanol to make ethanol unfit for human
consumption.
Ethanol from maize
Maize is crushed first and mixed with hot water. The mixture is then sieved and starch
(dextrin, a polymer of glucose) collects at the bottom of the collecting tank. The starch is
dried first and heated with dilute hydrochloric acid in a process called hydrolysis where
starch molecules are broken down in the presence water to produce glucose. The glucose
undergoes fermentation with yeast as before.
The ethanol produced is concentrated by fractional distillation.
note the fermentation processes are batch processes.
Ethanol from ethene
Ethane is reacted with steam in the presence of phosphoric acid (H 3PO4) catalyst and a temperature
of 300oC.
The benefits of producing ethanol this way is that
1. The process is continuous
2. Only one product is produced making the process particularly efficient.
Uses of ethanol
-As a solvent for dyes and medicines
-As a beverage/drink
-In blended fuels e.g. petrol blend
-In methylated spirits
-medical purposes
Global warming
Global warming is the increase in temperature of Earth’s atmosphere due to trapping of heat by
greenhouse gases.
39
greenhouse effect is the trapping of heat from sun by greenhouse gases to regulate Earth temperature
so that not all heat is reradiated back to space. However, increased industrialization releases more
greenhouse gases to atmosphere, contributing to Global Warming.
Greenhouse gases
Carbondioxide
Methane
Nitrous oxides
Chloroflurocarbons CFCs
Water vapour
ozone
Causes of global warming
1. Burning fossil fuels like coal and oil produces carbon dioxide and nitrous oxide.
2. Cutting down forests (deforestation). Trees help to regulate the climate by absorbing CO2 from
the atmosphere. So when they are cut down, that beneficial effect is lost and the carbon stored
in the trees is released into the atmosphere, adding to the greenhouse effect.
3. Veld fires release gases such as carbon dioxide
4. Cars emit carbon dioxide and nitrous oxide via the exhaust
5. Large scale livestock farming. Cows and sheep produce large amounts of methane when they
digest their food.
6. Fertilisers containing nitrogen produce nitrous oxide emissions.
Gas
Water vapour
Source
Oceans, lakes, rivers, reservoirs. Humans have little impact
upon levels.
Carbon dioxide
Burning of fossil fuels, and forests
Methane
Much from break down of organic matter by bacteria (rice
paddy fields) cows, swamps marshes
Ozone
Naturally from some oxygen atoms. Ozone in the troposphere
is due to chemical reactions between sunlight and agents of
pollution.
Chlorofluorcarbons
Fridges and aerosols.
Nitrous oxide
Nitrate fertilisers, transport and power stations (combustion)
contribution
98%
50%
18%
8%
25%
6%
Group 1
The elements in group one are also called alkali earth metals.
Physical properties
 They are good conductors of electricity and heat.
 They are soft metals uch that they can be cut with a knife
 They are metals with low densities with lithium, sodium and potassium having so low densities
that they can float on water.
 They have shiny surfaces when freshly cut with a knife however they quickly tarnish.
 They have low melting and boiling points. For example, lithium has a melting point of 181 °C
and potassium has a melting point of 64 °C.
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Chemical properties
They are all very reactive metals and they are stored under oil to prevent them coming into contact
with water or air. The reactivity increases as you go down the group.

They burn in oxygen or air, with characteristic flame colours, to form white solid oxides. For
example, lithium reacts with oxygen in air to form white lithium oxide, according to the
following equation:
lithium + oxygen →● lithium oxide
4Li(s) + O2(g) → 2Li2O(s)

They react vigorously with water to give an alkaline solution of the metal hydroxide as well as
producing hydrogen gas. For example:
potassium + water →● potassium hydroxide+ hydrogen gas
2K(s)

+ 2H2O(l) → 2KOH(aq) + H2(g)
These Group I oxides all dissolve in water to form alkaline solutions of the metal hydroxide.
lithium oxide + water →● lithium hydroxide
Li2O(s)+ H2O(l) → 2LiOH(aq)
Group II
the elements in group II are also called alkali earth metals because they are commonly found
naturally in minerals.
Physical properties
 They have low melting and boiling points but higher than than group I elements
 They have low densities but higher than group I elements
 They are harder than those in Group I.
 They are silvery-grey in colour when pure and clean. They tarnish quickly, however, when left in
air due to the formation of a metal oxide on their surfaces
 They are good conductors of heat and electricity.
Chemical properties
The elements are reactive but less reactive than group I elements.
The reactivity increases as you go down the group.
 They burn in oxygen or air with characteristic flame colours to form solid white oxides. For
example:
magnesium + oxygen → magnesium oxide

2Mg(s) + O2(g) → 2MgO(s)
They react with water, but they do so much less vigorously than the elements in Group I. they
react with water in varying degrees. Calcium readily reacts with cold water to form calcium
hydroxide
calcium + water → calcium hydroxide + hydrogen
Ca(s) + 2H2O(l) → Ca(OH)2(aq) +H2(g)
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Magnesium reacts extremely slow with cold water but readily with steam.
Group VII – the halogens
Group VII elements are all non metals. They are also called halogens.
Physical properties
 These elements are coloured and darken going down the group (Table 9.4).
Table 9.4 Colours of some halogens.
Halogen
Colour
Chlorine
Yellowish-green
Bromine
Red–brown
Iodine
black

They exist as diatomic molecules, for example Cl2, Br2 and I2.

They show a gradual change from a gas (Cl2), through a liquid (Br2), to a solid (I2) as the density
increases.
The boiling and melting points of the halogens increase as you move down the group.
Chemical properties
They form molecular compounds with other nonmetallic elements, for example HCl.
They react with hydrogen to produce the hydrogen halides, which dissolve in water to form
acidic solutions.



Uses of halogen


Bleaching agents- chlorine is used to bleach wood pulp in the paper making industry
Sterilisation- chlorine is used to kill bacteria and viruses in drinking water. Iodine is used in
medicine and disinfectants. Bromine is used to make disinfectants.

Manufacturing of other chemicals. Fluorine is used in the form of fluorides in drinking water
and toothpaste because it reduces tooth decay by hardening the enamel on teeth. Chlorine
is used to make PVC plastic.
Group 0
The group is also known as group VIII (8). The elements are called noble gases.
They are colourless, monoatomic gases.
They are very unreactive (inert).
Reactivity series
With air/oxygen
Many metals react directly with oxygen to form oxides. For example, magnesium burns brightly in
air/oxygen to form the white powder magnesium oxide.
magnesium + oxygen → magnesium oxide
2Mg(s) + O2(g) → 2MgO(s)
With water/steam
Reactive metals such as potassium, sodium and calcium react with cold water to produce the metal
hydroxide and hydrogen gas. Hydrogen makes a large popping sound with a lighted splint (test for
hydrogen). For example, the reaction of sodium with water produces sodium hydroxide and hydrogen.
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sodium + water → sodium hydroxide + hydrogen
2Na(s) + 2H2O(l) → 2NaOH(aq) + H2(g)
Magnesium reacts slowly with cold water but vigorously with steam to produce magnesium oxide
and hydrogen
magnesium + steam → magnesium oxide + hydrogen
Mg(s) + H2O(g) → MgO(s) + H2(g)
Zinc gives an oxide that is yellow when and white when cold
Reaction of Metals with Dilute Hydrochloric Acid
Pottasium, Sodium, Calcium, Magnesium, Zinc and Iron reacts with dilute acids. For example
Magnesium + hydrochloric acid → Magnesium Chloride + Hydrogen
Mg(s) + 2HCl(aq) → MgCl2(aq) + H2(g)
Lead reacts with concentrated hydrochloric acid.
Copper and Gold have no reaction with dilute nor concentrated hydrochloric acid
Summary of reactions
Reactivity series
Reaction with
air/oxygen
Reaction with
water
Reaction with
dilute acid
Potassium (K)
Burn very brightly
and vigorously
Produce H2 with
decreasing vigour
Sodium (Na)
Burn to form an
oxide with
decreasing vigour
Produce H2 with
decreasing vigour
with cold water
Calcium (Ca)
Magnesium (Mg)
React with steam
with decreasing
vigour
Aluminium (Al*)
Zinc (Zn)
Iron (Fe)
React slowly to
form the oxide
Lead (Pb)
Copper (Cu)
very slow oxidation
Silver (Ag)
Do not react
Do not react with
cold water or
steam
Do not react with
dilute acids
Gold (Au)
Platinum (Pt)
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The reactivity series
Potassium (K)
Sodium (Na)
Calcium (Ca)
Magnesium (Mg)
Aluminium (Al*)
[Carbon (C)]
Zinc (Zn)
Iron (Fe)
Lead (Pb)
[Hydrogen (H)]
Copper (Cu)
Silver (Ag)
Gold (Au)
Platinum (Pt)
decreasing reactivity
Displacement Reactions
Displacement reaction is the displacement of ions of metal from compounds of metals lower in
reactivity series by metals higher in reactivity series.
E.g. Magnesium displaces copper(II) chloride
Mg(s) + CuCl2(aq) → MgCl2(aq) + Cu(s)
But when Mg is placed in KCl, no reaction occurs.
Mg(s) + KCl2(aq) → No reaction
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