Download N5 Chemistry Summary notes 2017

Reaction Rates
Following the rate of a reaction.
Reaction rate is the speed of a reaction. The faster the reaction the higher the rate.
The progress of a chemical reaction can be followed by examining the reaction rate. There are several
methods that can be used to follow a reaction rate.


Change in volume of any gases produced
Change in mass of reaction mixture if gases are given off or taken in.
Other changes which can be monitored are: colour, pH, concentration and many more
As these changes happen over a period of time, the reaction rate is expressed as
Rate = Change in something
Change in time.
Changes in Volume
The volume of the gas can be measured using the following experiments.
gas
Measuring
cylinder
50cm3 of
1 mol/l
HCl
0.1 g of Mg
water
OR
Gas syringe
50cm3 of
1 mol/l
HCl
0.1 g of Mg
1
As the experiment proceeds the volume of gas is measured and from these results a graph can be drawn.
Time (min)
0 2 4 6 8
10 12 14 16
Volume of gas (g) 0 40 70 90 100 105 105 105 105
The volume increases as more gas is formed.
Volume of gas
produced (cm3)
Time (min)
The graph is steeper at the beginning as the rate is higher (i.e. the gas is made more quickly)
The graph levels off at 10 minutes when no more gas is being made (i.e. the reaction is finished)
In this experiment the rate would be calculated by:
Rate = Change in volume of gas (cm3)
Change in time (min)
And would have the units:
cm3/min
i.e. in the first 2 minutes
between 2 and 6 min
Rate = 40 - 0 – 0 (cm3))
2 – 0 (min)
Rate =90 – 40 (cm3)
6 – 2 (min)
= 40 (cm3)) / 2 (min)
= 50 (cm3) / 4 (min)
=20 cm3/min
= 12.5 cm3/min
2
Changes in mass
The mass change can be followed using the experiment below.
50cm3 of
1 mol/l
HCl
0.1 g of Mg
Balance
127.300g
As the experiment proceeds the mass is measured.
Time (min)
0
2
4
6
8
10
12
14
16
Mass of flask (g) 127.3 126.5 125.9 125.5 125.3 125.2 125.2 125.2 125.2
The mass falls as the gas is made and escapes.
From these results a graph can be drawn of mass of flask against time.
Mass of flask
(g)
Time (min)
Rate = Change in mass of flask (g)
Change in time (min)
3
And would have the units:
g/min
The mass of gas formed can also be measured by taking the mass at any time away from the initial mass at
time zero.
i.e.
Before the reaction starts (time zero) the mass
At 2 minutes the mass
Therefore the mass of gas formed
=
127.3 g – 126.5g
=
=I
=
1277.3
126.5g
0.8g
We can therefore work out another row for our table: ‘Mass of gas formed’.
Time (min)
0
2
4
6
8
10
12
14
16
Mass of flask (g)
127.3 126.5 125.9 125.5 125.3 125.4 125.4 125.4 125.4
Mass of gas formed (g) 0.0
0.8
1.4
1.8
2.0
2.1
2.1
2.1
2.1
This will give a different shaped graph.
From these results a graph can be drawn of mass of flask against time.
Mass of gas formed
(g)
Time (min)
Rate = Change in mass of gas formed (g)
Change in time (min)
would have the units: g/min
The rate can be calculated from the graph or table
Rate = Change in mass of gas formed (g)
Change in time (min)
i.e. in the first 2 minutes
between 2 and 6 min
Rate = 0.8 – 0.0 (g)
2 – 0 (min)
Rate =1.8 – 0.8 (g)
6 – 2 (min)
= 0.8 (g) / 2 (min)
= 1.0 (g) / 4 (min)
= 0.4 g/min
= 0.25 g min
The rate is faster at the start of the reaction, which we can tell from the graph because it is steeper.
4
And
Atomic structure
Everything is made up of tiny particles called atoms.
Atoms are mostly empty space made up of smaller sub-atomic particles.
At the centre of the atom is the nucleus.
This contains two types of particles, called protons
and neutrons.
Spinning around the nucleus are very fast
moving particles called electrons.
They move in different levels, called shells.
Protons have a positive electrical charge.
Neutrons have no electrical charge and are neutral.
Moving around the nucleus are very small negatively charged electrons.
Nucleus containing the Protons and Neutrons
(positive)
(positive)
(neutral)
Electron Cloud containing the Electrons
(negative)
(negative)
Name of particle
Proton
Neutron
Electron
Electric Charge
Mass (amu)
1
1
1/1840 (almost zero)
1+
0
1-
Position
Nucleus
Nucleus
Electron Cloud
The number of protons in the nucleus of an atom is called the Atomic number. Every
element has a different atomic number.
The atom itself is neutral because the number of protons and electrons are the same.
Therefore the positive charge of the nucleus is equal to the sum of the negative charges of
the electrons.
Atomic Number = Number of protons
(also number of electrons)
5
Every atom has two numbers.
Atomic Number = number of protons
Mass Number
= number of protons + neutrons
Electrons are not counted when calculating the mass number as they add so little to the
atom’s mass.
Using a system called Nuclide notation, scientists can show the symbol, mass number and
atomic number of an element.
Examples:
7
3
Li
MassNumber
19
AtomicNumber
9
F
Electrons move around the nucleus in orbits. The electron orbits are called shells.
Each shell can hold only a certain number of electrons.
The first shell
The second and third shells
The last or outermost shell
can hold a maximum of 2 electrons.
can hold a maximum of 8 electrons.
can hold a maximum of 8 electrons.
Electrons further away from the nucleus have more energy then those close to the nucleus.
E.g. an atom of the element carbon has an atomic number of 6.
Its electrons are arranged so that each electron in the second shell
is in a separate quarter of the shell.
electron arrangement 2,4
6
An atom of the element oxygen has an atomic number of 8.
Its electron arrangement is 2,6
O
Electron arrangement and properties
Sodium has an atomic number of 11.
So its electron arrangement is 2,8,1
Na
If we look at Group 1 (the alkali metals)
These metals all react violently with water and form alkalis. They all have similar chemical
properties.
All of the alkali metals one electron in their outer shell.
Lithium
Sodium
Potassium
Rubidium
Caesium
Li
Na
K
Rb
Cs
2,1
2,8,1
2,8,8,1
2,8,18,8,1
2,8,18,18,8,1
If we look at Group 8 (0) (the Noble)
All the elements in group 0, the noble gases, are inactive and have virtually no chemistry.
The outer shell is Full, (8 outer electrons, except He which is full with 2 outer electrons).
Helium
Neon
Argon
Krypton
Xenon
Radon
He
Ne
Ar
Kr
Xe
Rn
2,
2,8
2,8,8,
2,8,18,8
2,8,18,18,8,
2,8,18,32,18,8,
Ne
It is the number of electrons in the outer shell, which gives atoms their properties.
7
Isotopes
All atoms of an element must contain the same number of protons, but they can contain different numbers of
neutrons.
We can therefore get atoms of the same element with different masses.
These are called isotopes.
Isotopes:
atoms with:
Same number of protons
but
different
number of neutrons
but
different
Mass Number
OR
Same Atomic Number
35
37
17
17
Cl
Cl
e.g.35Cl has 17 protons and 18 neutrons
37
Cl has 17 protons and 20 neutrons
RELATIVE ATOMIC MASS (RAM)
(page 4 of the data book)
Most elements contain a mixture of isotopes.
e.g.
37
35
exist in the proportions
25%
75%
Cl
Cl
The relative atomic mass is the average mass of all the isotopes an element has. It is rarely a
whole number but has been rounded off to the nearest 0.5.
Relative Atomic Masses in the data book as always closest to the most abundant isotope.
The ram for chlorine (on page 4 of the data book) is 35.5.
This is because there is more 35Cl than 37Cl
It is possible to get equal proportions isotopes.
107
109
e.g.
Ag
Ag
50%
50%
The ram is therefore 108 (i.e. half way between 107 and 109).
8
Covalent Bonding
Non-metal elements
Group 0 elements, the noble gases exist as single atoms, the monatomic gases.
These elements are very unreactive.
The rest of the non-metallic elements exist as molecules which are groups of atoms joined
together with covalent bonds.
Some elements exist as diatomic molecules (a molecule made of two atoms)
There are seven non-metal elements in the periodic table which occur as diatomic
molecules.
They are : hydrogen (H2)
oxygen (O2)
O
nitrogen (N2)
fluorine (F2)
chlorine (Cl2)
bromine (Br2)
iodine
(I2)
Compounds
Covalent compounds
Covalent compounds contain only non-metal elements
eg hydrogen oxide (pure water).
Both hydrogen and oxygen are non-metals,
so hydrogen oxide is covalent.
9
Covalent Bonding (non-metal/non-metal)
The Noble gases are stable elements as they have a full outer electron shell.
Other elements react until their atoms obtain a full outer shell and become stable.
Non-metal atoms obtain a full outer shell by sharing their outer electrons with other nonmetal atoms.
The sharing of outer electrons is called a COVALENT BOND.
The substance formed is then called a covalent molecule.
For example chlorine atoms have 7 outer electrons by two atoms sharing one electron each
they both obtain a full outer shell.
Cl
Chlorine atom
(7 outer electrons)
Cl
Cl
Chlorine atom
Cl
Chlorine molecule (Cl 2)
(7 outer electrons)
(each atom now has 8 outer electrons)
One shared pair of electrons is called a covalent bond.
A full structural formula for a Cl 2 molecule is drawn as
Cl
Cl
It is possible to form double and triple covalent bonds.
Double bonds
O
Oxygen atom
(6 outer electrons)
O
O
Oxygen atom
O
Oxygen molecule (O 2)
(6 outer electrons)
(each atom now has 8 outer electrons)
The full structural formula for the oxygen molecule O 2 is O O
10
Triple Bonds
N
N
Nitrogen atom
Nitrogen atom
(5 outer electrons)
(5 outer electrons)
N
N
Nitrogen molecule (N 2)
(each atom now has 8 outer electrons)
The full structural formula for the nitrogen molecule N 2 is N
This contain triple covalent bond.
11
N
Compounds.
We can apply the same rules to covalent compounds.
Covalent bond
Hydrogen chloride
H
H
Cl
+
Hydrogen atom
1 outer electron
Chlorine atom
7 outer electrons
Cl
Hydrogen chloride
molecule (HCl)
Each atom in a hydrogen chloride molecule has a full outer shell.
The hydrogen atom has 2 outer electrons like the noble gas helium.
The chlorine atom has 8 outer electrons like the noble gas argon.
Covalent molecules can be made from more than two atoms.
Water (H2O)
H
O
Hydrogen has 1 outer electron
It needs to obtain 1 more electron
Oxygen has 6 outer electrons
It needs to obtain 2 more electrons
The solution is for the oxygen atom to share its electrons with 2 hydrogen atoms.
O
H
H
The full structural formula for H2O is
O
H
12
H
Covalent Bonds
Covalent bonds hold the atoms in a molecule together.
Take Hydrogen (H2) as an example.
Each hydrogen atom has 1 positive proton in the nucleus and 1 negative electron in its
electron shell.
Two hydrogen atoms share their outer electron to obtain a full outer shell.
+
+
-
The negatively charged electrons are attracted to both positively charged nuclei and this
holds the atoms together.
13
Formulae
The number of electrons which an atom needs is called its valency.
This tells us the number of covalent bonds that it will form.
Using this we can work out the full structural formula
For example: Water
Hydrogen has 1 outer electron; it needs 1 more outer electron and so forms 1 bond.
Oxygen has 6 outer electrons; it needs 2 more outer electrons and so forms 2 bonds.
We therefore draw the full structural formula
O
H
H
As all elements in the same group of the periodic table have the same number of outer
electrons we can save time by looking at the group
Group
Group 4 Group 5 Group 6 Group7
Number of outer electrons
4
5
6
7
Valency (Number of electrons needed )
4
3
2
1
Number of bonds formed
4
3
2
1
For example Phosphorus chloride
Element
Phosphorus (P) Chlorine (Cl)
Group
5
7
Number of outer electrons
5
7
Valency (Number of electrons needed )
3
1
Number of bonds formed
3
1
Full structural formula
Cl
P
Cl
Cl
For example carbon sulphide
Element
Carbon (C) Sulphur (S)
Group
4
6
Number of outer electrons
4
6
Valency (Number of electrons needed )
4
2
Number of bonds formed
4
2
Full structural formula
S
C
14
S
Working out molecular formulae
Molecular formulae use symbols to show the number of each atom in a molecule. We can
deduce the molecular formula from the structural formula or work it out using the valencies.
Example Phosphorus Chloride
Symbols
Valency
Swap
Formula
P
3
1
P1
Cl
1
3
Cl3
The formula for phosphorus chloride is PCl3
Sometimes you will need to simplify the formulae.
Example Carbon sulphide
Symbols
Valency
Swap
Simplify
Formula
C
4
2
1
C1
S
2
4
2
S2
The formula for carbon sulphide is CS2
As Carbon sulphide would work out as C2S4
This is simplified to CS2 to get the correct number of atoms as show by the full structural
formula;
S C S
15
Compounds which do not obey the rules.
Some covalent compounds do not obey the valency rules.
The names of these compounds tell us the molecular formulae.
eg
In the compound carbon dioxide, the prefix di means two.
So the compound contains one carbon and two oxygen atoms.
The molecular formula is CO2
Prefixes in compound names are:
mono
di
tri
tetra
penta
=
=
=
=
=
1
2
3
4
5
Examples of compounds where the name tells us the formula are:
carbon monoxide
CO
nitrogen dioxide
NO2
sulphur trioxide
SO3
dinitrogen tetraoxide
N2O4
(note in these formulae we do not simplify)
16
Shapes of covalent molecules
The electrons in the first electron shell are found in a sphere but the other electron shells
actually hold the electrons in four orbitals.
Thus hydrogen and oxygen which we represent as;
H
O
Should really be represented as
and
This would give us the molecule water drawn as:
We do not have to use orbitals when we show how molecules form, however it is important
when working out the actual shape of molecules
Each orbital can hold two electrons giving the 8 outer electrons which atoms need to have a
stable outer shell.
17
When molecules form, the orbital’s overlap and gives us one of four possible shapes.
The shape depends on the number of bonds formed between the atoms.
Number of Bonds
Shape
Name of shape
1
H
F
Linear
2
Angular
3
Pyramidal
4
Tetrahedral
18
Covalent Molecules and Covalent Networks
Covalent substances can be divided into two groups depending on their structure.
Covalent molecules
So far we have only considered covalent molecules where the molecular formula tells us the
exact number of atoms in a molecule.
eg water
H2O
A molecule of water has 1 oxygen atom
joined to two hydrogen atoms.
Covalent networks
A few covalent substances have their atoms joined into a huge network.
e.g. diamond (carbon) (C)
and
silicon dioxide (SiO2)
These substances will contain many millions of atoms.
Their formulae do not tell us the number of each atom but the ratio of the atoms.
e.g. The formula for silicon dioxide is SiO2.
This does not mean that the silicon dioxide network contains 1 silicon atom joined to 2
oxygen atoms.
It means the ratio of Si : O is 1:2
i.e. there are twice as many oxygen atoms as silicon atoms.
19
Ions
Atoms of the same element always have the number of protons but the number of electrons
can change when a compound is formed. This gives the atom a charge and we call it an ion.
Metal atoms form positive ions
Non-metal atoms form negative ions.
Positive and negative ions are found together in some compounds.
A positive ion is made when an atom loses electrons
A negative ion is made when an atom gains electrons.
For example:
Atom Symbol
Na
Electrons Lost or Gained
1 electron lost
Ion formed
Na
Al
3 electrons lost
Al
Cl
1 electron gained
Cl
O
2 electrons gained
O
What information do we get from an ion symbol?
Here is a symbol for the magnesium atom.
Mass number
25
Atomic number
12
Mg
This magnesium atom contains :- 12 protons
13 neutrons
12 electrons
Here is a symbol for the magnesium ion.
Mass number
25
Atomic number
12
magnesium ion contains :-
Mg 2+
Electric charge (for ions only)
12 protons
13 neutrons
10 electrons
Note: in an ion the number of protons is not equal to the number of electrons.
20
+
3+
2-
Ionic Bonding (metal / non-metal)
In order to become stable elements want the stable electron arrangement as their nearest
Noble gas (8 Outer electrons, 2 for Li & Be)
Elements can also obtain the stable electron arrangement of the noble gases by forming
charged particles called IONS.
Metal atoms form positive ions by losing electrons.
loses 1 e
sodium atom
sodium ion
Na+
Na
Electron
arrangement
Na+
Na+
Na
2,8,1
+ e-
2,8,
number of positive charges = number of electrons lost
= periodic table group number
Non-metal atoms form negative ions by gaining electrons.
Cl -
Cl
gains 1 e
e.g.
chlorine atom
Cl
Electron
arrangement
+e
chloride ion
-
Cl-
2,8,7
2,8,8
number of
= number of
negative charges
electrons gained
These types of equations are known as an ion-electron equations and can be
written for the formation of any simple ion.
21
For example
Magnesium atom loses 2 e to form a Magnesium ion
Mg
2,8,8,2
Mg 2 ++
2,8,8,
2e
and
Sulphur atom gains 2e to form a Sulphide ion
S + 2e
2,6
S22,8,
Note : The names of the non-metals change to have the ending ‘ –ide ‘ when an ion is
formed.
Transition metal ions
The charge on the metal ion is shown by a Roman number after the name,
e.g.
an iron (III) ion is Fe3+
a copper (II) ion is Cu2+
Complex ions
Some atoms form groups of atoms which have an overall charge. These are listed on page 4
of the data book.
e.g. sulphate
SO42nitrate
NO3 The names of these ions usually end in ‘ ate ‘ or ‘ ite ‘ and contain oxygen.
22
Ionic Compounds (Metal/ Non-metal)
We know:
metals become stable by
Non-metals become stable by
losing electrons
gaining electrons
Ionic compounds are formed when metals transfer their outer electrons to the non-metals.
Eg Sodium wants to lose 1 electron and Chlorine wants to gain 1 electron
Na atom
2,8,1
Cl atom
2,8,7
1 electron transferred
Na+ ion
2,8,
Cl- ion
2,8,8,
Sodium chloride Na+ Cl-
Eg Calcium wants to lose 2 electrons and Chlorine wants to gain 1 electron
Ca atom
Cl atom
Cl atom
2,8,8,2
2,8,7
2,8,7
1 electron transferred
1 electron transferred
Ca2+ ion
2,8,8,
Cl- ion
2,8,8,
Cl- ion
2,8,8,
Calcium chloride Ca2+ (Cl-)2
23
Formulae of ionic compounds
e.g.
aluminium oxide
iron (III) bromide
a)
Al
O
Fe
Br
V
S
S
F
3
2
2
2
3
3
3
1
1
1
3
3
Al2O3
copper (II) nitrate
Cu
NO3
2
1
1
FeBr3
1
2
2
Cu(NO3)2
Note;
i)
The valency is obtained from the group in the periodic table;
ii)
iv)
1
2
3
4
5
6
7
8
Valency
1
2
3
4
3
2
1
0
metals give positive ions.
non-metals give negative ions.
therefore;
iii)
Group
Group
1
2
3
4
5
6
7
8
Valency
1+
2+
3+
4+/-
3-
2-
1-
0
The valency of transition metals is given in roman numerals.
If it is not given Assume:
silver
All other transition metals
Complex ions are obtained from the data book page 4.
Complex ions are identified by the name not ending in –ide.
The exceptions are hydroxide (OH-) and ammonium (NH4+).
24
valency (I)
valency (II)
STRUCTURE OF COMPOUNDS
Ionic compounds form a neat array oppositely charged ions.
e.g. sodium chloride Na+ Cl-
This structure is called an Ionic lattice (sometimes a crystal lattice).
The attractions between oppositely charged ions are very strong and difficult to break..
An Ionic formula does not tell use the actual number of ions present.
It tells us the ratio of ions
i.e.
Na+ Cl-
is
1:1
for every one Na+ ion there is one Cl- ion.
25
Acids and Bases
The pH Scale
 A continuous scale from below 0 to above 14
 It is a measure of the hydrogen ion (H+) concentration of a solution
Dissociation of Water
 Water exists as an equilibrium between water molecules and hydrogen and hydroxide
ions
H2O (l)
molecules
⇌H
+
(aq) + OH – (aq)
ions
 In water equilibrium lies over to the left – mainly exists as
In water and other neutral solutions (pH 7)
the concentration of H+ = the concentration of OH-
26
pH
OH- versus H+
Examples
concentration of
Hydrochloric acid,
Nitric acid,
Sulphuric acid
pH less than 7
Acidic
-
H+  OH
concentration of
Neutral
pH = 7
Alkaline
pH more than 7
-
H+ = OH
concentration of
-
+
H < OH
Water,
Sugar solution,
salt solution
Sodium hydroxide,
Potassium hydroxide
Bases/Alkalis
A base is a substance that neutralises and acid (brings the pH of the acid up to 7 ).
An alkali is a base which is soluble in water.
Making Acids or Alkalis
 Acids
Acids are made by dissolving soluble non-metal oxides in water.
o These make acidic solutions by raising the H+ concentration of water
o CO2 (forms carbonic acid, H2CO3 when dissolved in water)
o NO2 (forms nitric acid, HNO3 when dissolved in water)
 Alkalis
Alkalis can be made by dissolving soluble metal oxides in water to make metal
hydroxides - (see data book). For example:
Na2O + H2O
2 NaOH
-
o This makes an alkaline solution because it raises the OH concentration of the
water
o Common bases are metal oxides, metal hydroxides or metal carbonates. For
example:
 Na2O
 Na2OH
 Na2CO3
 Neutral
Insoluble oxides will not affect the pH of water.
They will not lower or raise the pH of water from 7 because they do not dissolve into
the water.
27
Common acids and alkalis used in the laboratory
hydrochloric
nitric
sulphuric
ethanoic
Common Laboratory Acids
HCl
HNO3
H2SO4
CH3COOH
Common Laboratory Bases
sodium hydroxide
NaOH
Potassium hydroxide
KOH
calcium hydroxide
Ca(OH)2
ammonia
NH3
Dilution of acids and alkalis
Acids


Acids have a pH below 7, which means that they have more H+ than OH-.
Diluting an acid, increases the pH towards 7 (concentration of H+ decreases)
Alkalis


-
Alkalis have above 7, which means that they have more OH than H+.
Diluting an alkali, decreases the pH towards 7 (concentration of OH- decreases)
28
Neutralisation (Salt Preparation)
Neutralisation is the reaction of acids with bases in which salt and water are produced.
 Base – a substance that reacts with an acid to neutralise it
eg metal oxide, metal hydroxide or metal carbonate
 Salt – a substance in which the hydrogen ion of an acid has been replaced by a metal or
ammonium ion
During a neutralisation the pH of the resulting solution is pH 7
This can easily been seen by using an indicator.
Neutralisation Reaction Equations
Acid
+
Base
Salt + Water
Acid
+
metal oxide
salt
+ H2O
Acid
+
metal hydroxide
(alkali)
salt
+ H2O
Acid
+
metal carbonate
salt
+ H2O + CO2
29
Naming of Salts
 The name metal in the base appears in the name of the salt
eg copper (II) carbonate produces a copper (II) salt
 The name of the negative ion in the acid appears in the name of the salt
acid
Hydrochloric
Nitric
Sulphuric
negative ion
Chloride
Nitrate
Sulphate
name of salt
Chloride
Nitrate
Sulphate
Examples:
nitric acid + iron (II) oxide
HNO3
+
FeO
hydrochloric + copper (II)
hydroxide
chloride
iron (II) nitrate + water
Fe(NO3)2
+ H2O
copper (II)
+ water
+ H2O
acid
HCl
+ Cu(OH)2
CuCl2
sulphuric
carbonate
+ calcium
sulphate
calcium + water + carbon acid
dioxide
H2SO4
+ CaCO3
CaSO4 + H2O + CO2
30
Neutralisation reactions
In neutralisation reactions the hydrogen ions from the acid are reacting with ions from the
base to form water molecules.
This can be seen by omitting the spectator ions, (the ions which do not change
during the reaction).
For example when hydrochloric acid is neutralised by sodium hydroxide we can see what is
happening by following 4 steps.
1. Write a balanced equation including state symbols
2HCl(aq) + NaOH(aq)
NaCl(aq) + H2O (l)
2. Rewrite including ionic formulae for underlined substances
-
-
H+Cl (aq) + Na+OH
-
Na+Cl (aq) + H2O(l)
3. Identify the ions which do not change -Spectator Ions
-
-
-
H+Cl (aq) + Na+OH
Na+Cl (aq) + H2O(l)
(sodium and chloride ions are the spectator ions in this equation)
4. Rewrite the equation missing out the spectator ions
H+ (aq)
+
OH
-
H2O(l)
31
Titrations
A titration is a method of analysis that allows you to determine the precise endpoint of a
reaction and therefore the precise quantity of reactant in the titration flask.
Neutralisation reactions can be carried out very accurately by using a burette and a pipette in
a titration. The volumes obtained can be used to calculate the concentrations of the acid or
alkali.
Pipette Filler
Pipette
Conical flask
32
FORMULAE
When determining formula we use valency which we can obtain from the group in the
periodic table.
1
1
1+
2
2
2+
3
3
3+
4
4
4-
H
Li
Na
K
Rb
Cs
Fr
Be
Mg
Ca
Sr
Ba
Ra
B
Al
Ti V Cr Mn Fe Co Ni Cu Zn Ga
Zr Nb Mo Tc Ru Rh Pd Ag Cd In
Hf Ta W Re Os Ir Pt Au Hg Tl
C
Si
Ge
Sn
Pb
Sc
Y
La
Ac
5
6
5
6
3 - 2-
N
P
As
Sb
Bi
O
S
Se
Te
Po
7
7
1-
8/0 Group
0 Valency
0 Charge
F
Cl
Br
I
At
He
Ne
Ar
Kr
Xe
Rn
Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu
Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No Lr
To determine a formula cross over the valences and simplify
e.g.
symbol
nitrogen fluoride
N
F
valency
3
1
symbol
carbon
C
oxide
O
valency
4
2
Formula
N1F3
Formula
C2O4
Simplify
NF3
Simplify
CO2
Ionic Formulae
In ionic formulae we show the charges.
Note
 The valency of transition metals is given in roman numerals.
 Complex ions are obtained from the data book page 8.
e.g.
a)
b)
calcium chloride
Ca2+
ClCa2+
(Cl-)2
Formulae: Ca2+(Cl-)2
iron (III) bromide
Fe3+
BrFe3+
(Br-)3
Fe3+(Br-)3
33
copper (II) nitrate
Cu2+
NO3Cu2+
(NO3-)2
Cu2+(NO3-)2
What does a formula tell you?
What a formula tells you depends on the structure of a substance.
Covalent molecules
The formula gives the number of atoms present in the molecule>
e.g. H2O contains 2 H atoms and 1 O atom
Covalent Network
The formula gives the simplest ratio of atoms in the substance.
e.g. SiO2 contains Si and O atoms in the ratio 1:2
Ionic Lattice
The formula gives the simplest ratio of ions in the substance.
e.g. Fe3+(Br-)3 contains Fe3+ and Br- ions in the ratio 1:3
34
Balanced equations
In a reactions atoms / ions are not destroyed they just ‘rearrange’ what other atoms / ions
they are attached to.
In an equation there must be the same number of each atom/ion on each side of the equation.
For example when magnesium reacts with oxygen gas to form magnesium oxide.
Mg
+
O2
MgO
There must be two oxygen on the right. For this to be possible we must have 2MgO
Mg
+
O2
2MgO
For this to be possible we must have started with 2 Mg
2Mg
+
O2
2MgO
+
O2
2MgO
Thus the balanced equation is:
2Mg
We describe this as 2 moles of magnesium reacting with 1 mole of oxygen forming 2 molesa
of magnesium oxide.
2Mg
2 mol
+
O2
1mol
2MgO
2 mol
35
Mole calculations
Using the data booklet page 7 we can work out the gram formula masses (GFM) of elements
and compounds.
GFM is defined as the mass f 1 mole.
Calcium nitrate
Ca(NO3)2
6 x O = 6 x 16 = 96
2 x N = 2 x 14 = 28
1 x Ca = 1 x 40 = 40
GFM = 144g = 1 mole
Once we know the GFM we can calculate masses and number of moles using the triangle
m
Number of moles
n
Mass
x GFM
What is the mass of 0.1 moles of calcium nitrate?
m
=
n x GFM
=
0.1 x 144g
=
14.4g
How many moles of calcium nitrate are there in 72g of calcium nitrate?
n
=
=
=
m/GFM
72/144
0.5 moles
36
Calculations from balanced equations
To do this use the following technique:
Here are the steps to carry out:
1. Put a ? over what you need to calculate.
2. Put a  over what you know about.
3. Use the balanced equation to write mole ratio.
4. Calculate the mass of each substance
5. Using what you actually have calculate what you want to know
What mass of MgO will be formed when 98g of Mg is burned in excess oxygen?
√
2Mg
+
Mole ratio 2 mol
2(24.5)
49g
Actually have 98g
?
2MgO
2 mol
2(24.5 + 16)
81g
98 x 81
49
= 162g
O2
37
Concentration of Solutions
Concentrations are expressed as the number of moles of a substance dissolved in 1 litre of
water. (mol l-1).
We use the triangle
Once we know the GFM we can calculate masses and number of moles using the triangle
n
concentration
(mol l-1)
number of moles
c x V
volume (in litres)
What is the concentration if 0.1 moles of calcium nitrate is dissolved in 250cm3 of water?
c
=
n/V
250 cm3 = 250/1000 l = 0.25 l
=
0.1 / 0.25
=
0.4 mol l-1
What mass of Calcium nitrate is needed to make 100cm3 of a 0.1 mol l-1 solution?
m
=
=
=
=
=
n x GFM
n x 144g
n x 144g
0.01 x 144
1.44g
We work out GFM from the formula as shown above.
To calculate n we use: n = c x v
= 0.1 x v
v = 100cm3
= 100/1000 l
= 0.1 l
= 0.1 x 0.1
= 0.01 mol
38
Calculations for titrations
To do this use the following technique:
Here are the steps to carry out:
1. Put a ? over what you need to calculate.
2. Put a  over what you know about.
3. Use the balanced equation to write mole ratio.
4. Calculate the number of moles you actually have of the substance you know about
5. Use the mole ratio to get the number of moles of the substance asked about.
6. Calculate the answer to the question
Remember to turn all volume into litre by dividing cm3 by 1000.
i.e. 20cm3 = 20/1000 = 0.020 litres
Example
If 20cm3 of NaOH is neutralised by 20cm3 of 1 mol l-1 H2SO4, what is the concentration of
the NaOH?
Mole ratio
√
H2SO4
1 mol
+
?
2NaOH
2 mol
Na2SO4+2H2O
Actually have
c =1 mol l-1
v = 20cm3 = 0.020 litre
n = 1 x 0.02
= 0.02 mol
v = 20cm3 = 0.020 litre
c=?
c = n/v
0.04 mol
C = n/v
= 0.04 / 0.02
= 2 mol l-1
39
Unit 2
Hydrocarbons.
Burning hydrocarbons always produces carbon dioxide and water.
Results
U-tube
Colourless liquid condensed
Cobalt chloride paper
changes from blue to pink.
Water present
Conclusion
Boiling tube
Lime water turns cloudy
Carbon dioxide present
This tells us that hydrocarbons contain
Carbon : because carbon dioxide is made
Hydrogen: because water is made
It does not tell us that they contain oxygen because the oxygen in carbon dioxide and
water can have come from the air on burning.
Word Equation :
methane
+
oxygen
→
carbon dioxide + water
Formula equation :
CH4
+
2O2
→
CO2
40
+ 2H2O
Measuring the Energy
The energy produced when a fuel is burned is calculated using the equation:
Eh = cmΔT
where
Eh = energy gained by the water
c = specific heat capacity of water = 4.18 kJ0C-1kg-1
m = mass of water (1 litre = 1kg)
ΔT= change in temperature of water
For example when 200cm3 of water is heated from 20oC to 50oC using the following
apparatus.
c = specific heat capacity of water = 4.18 kJ0C-1kg-1
m = 200 / 1000 = 0.2kg
ΔT= 50 – 20 = 30oC
Eh
= 4.18 x 0.2 x 30
= 25.08 kJ
As the reaction is exothermic we write:
Eh
= - 25.08kJ
41
Hydrocarbons

Hydrocarbons are molecules containing hydrogen and carbon only
Carbon
Hydrogen-
–
-
always forms 4 bonds
always forms 1bond
The hydrocarbons are grouped into families called Homologous series.
Homologous Series- A group of hydrocarbons with:
 Similar chemical properties
 Same General Formula
42
Alkanes
General formula CnH 2n+2
Name
No. of Carbon atoms
Molecular formula
Methane
1
CH4
Ethane
2
C2H6
Propane
3
C3H8
Butane
4
C4H10
Pentane
5
C5H12
Hexane
6
C6H14
Heptane
7
C7H16
Octane
8
C8H18
Structural Formula
Methane
H
H
CH4
C
H
H
Ethane
H
Propane
H
H
H
C2H6
C
C
H
H
H
H
H
C
C
C
H
H
H
H
C3H8
H
Note: the bonds between the atoms are covalent single bonds
43
Alkenes

General Formula CnH2n
Ethene
H
Propene
H
H
H
C2H4
C
C
H
H
H
C
C
C
H
C3H6
H
H
Note: Alkenes contain a double bond (C=C).
Alkene
Ethene
Propene
Butene
Pentene
Hexene
No. of Carbon Atoms
Molecular Formula
2
C2H4
3
C3H6
4
C4H8
5
C5H10
6
C6H12
44
Cyclo alkanes

General Formula CnH 2n
Cyclopropane
H
H
C
H
C
C
H
H
Cyclobutane
C3H6
H
H
H
C4H8
H
C
C
H
H
C
C
H
H
H
Alkane
Number of Carbon atoms
Molecular formula
Cyclopropane
Cyclobutane
Cyclopentane
Cyclohexane
3
4
5
6
C3H6
C4H8
C5H10
C6H12
45
Saturated or unsaturated hydrocarbons.
Alkanes and cycloalkanes are saturated because all the (C-C) bonds are single bonds.
Alkenes are unsaturated because they contain a (C=C) double bond..
The double bond is responsible for the quick reaction with bromine water.
Bromine can be used to tell the difference between saturated and unsaturated hydrocarbons.
Addition Reactions
 The bromine adds to the alkene across the double bond.

The reaction involves the breaking of the carbon-carbon double bond to form a carbon-carbon single
bond.
C3H6
H
+
Br2
C3H6Br2
H
C
C
H
H
H
+
Br
Br
H
C
H
C
Br Br
Changing Alkenes to Alkane
Alkenes can react with Hydrogen under certain conditions to form Alkanes.
Ethene + Hydrogen  Ethane
C2H4 +
H2  C2H6
Propene + Hydrogen  Propane
C3H6
+
H2
 C3H8
46
H
Isomers
Isomers: same molecular formula but different structural formula.
Butane – formula is C4H10 and there are two different structures.
H
H
H
H
H
C
C
C
C
H
H
H
H
H
This is a straight chain molecule.
This is a branched chain molecule.
For the molecule C4H8
H
H
H
C
C
H
H
C
C
H
H
H
Cyclobutane
H
H
H
H
C
C
C
C
H
H
H
Butene
47
H
Naming Alkanes
Identify the longest chain.
Number the carbon atoms to place the branch on the lowest number possible.
Identify the position of the chains.
Identify the number of each branch.
2,3,4-trimethylpentane
Naming Alkenes
This is similar to alkanes but th position of the double bond must be as low as possible
.
3 – methybut – 1 – ene
48
Alkanols (alcohols)
General formula CnH2n+1OH
Alkanols contain the functional group –OH (hydroxyl)
Name
No. of Carbon atoms
Molecular formula
Methanol
1
CH3OH
Ethanol
2
C2H5OH
Propanol
3
C3H7OH
Butanol
4
C4H9OH
Pentanol
5
C5H11OH
Hexanol
6
C6H13OH
Heptanol
7
C7H15OH
Octanol
8
C8H17OH
Structural Formula
Methanol
H
H
CH3OH
C
OH
H
Ethanol
H
Propanol
H
H
H
C
C
H
H
H
H
C2H5OH
OH
(CH3CH2OH)
H
C
C
C
H
H
H
C3H7OH
OH
(CH3CH2CH2OH)
49
Naming alkanols
Alkanols are named like the alkenes but here the number of the
hydroxyl group (-OH) must be as low as possible.
Butan-1-ol
Butan-2-ol
50
Uses of Alkanols (alcohols)
Alcoholic Drinks.
Ethanol is found in alcoholic drinks.
Fuels
Alcohols burn very easily with a clean flame and so they can be used
as a fuel.
The products of combustion are carbon dioxide and water.
C2H5OH +
3O2
2CO2
+
3H2O
Ethanol can be used as a petrol replacement.
Methanol is used as the fuel for drag racers.
CH3OH
+
1½ O2
CO2 +
2H2O
Solvents
Alkanols are good solvents for many types of substances.
The carbon chain helps alkanols dissolve in covalent substances
The hydroxyl group (-OH) helps alkanols dissolve in water.
Ethanol is used as a solvent in perfumes, aftershave and mouthwash
as it is good at dissolving and evaporates easily as well.
51
Alkanoic acids (carboxylic acids)
General formula CnH2n+1COOH
Alkanoic acids contain the functional group –COOH (carboxyl)
O
C
OH
Name
No. of Carbon atoms
Molecular formula
Methanoic
1
HCOOH
Ethanoic
2
CH3COOH
Propanoic
3
C2H5COOH
Butanoic
4
C3H7COOH
Pentanoic
5
C4H9COOH
Hexanoic
6
C5H11COOH
Heptanoic
7
C6H13COOH
Octanoic
8
C7H15COOH
Be careful when writing the molecular formula or naming alkanoic acids. The stem will have
1 carbon less than you expect as it also contains the carboxyl group COOH.
i.e. C4H9COOH is pentanoic acid as there are 5 carbon atoms in this alkanoic acid.
52
Structural Formula
Methanoic acid
O
H
Ethaoic acid
H
HCOOH
C
OH
H
O
CH3COOH
C
C
OH
H
H
O
C
C
C
H
H
H
Propanoic acid
H
C2H5COOH
OH
(CH3CH2COOH)
Uses of Alkanoic acids (carboxylic acids)
 food preservatives
(i.e. vinegar is a solution of ethanoic acid used for pickled onions)
 cleaning
products and descaling agents
(i.e. vinegar in solutions for washing windows)
 manufacture
of esters and plastics
53
Esters
Esters are made by reacting alkanols with alkanoic acids.
Uses of Esters
Food flavourings
Esters are added to foods as their smell influences the flavour, for example
pineapple cubes contain methyl butanoate
Fragrances.
The smells of specific esters allows them to be used in things such as perfumes
and air fresheners.
Solvents.
Many esters are sold to remove stains from clothes and house furnishings.
Materials
Polyesters are esters used in clothing and fabrics such as nylon.
54
POLYMERS
Polymers are macromolecules (big molecules).

Monomers are the small molecules that combine to form a polymer.

A polymer is the large molecule formed by combining many monomers
Polymers are very large molecules containing many repeating units.
Plastics are made by a process called polymerisation of which there are two types:
Addition polymerisation
Condensation polymerisation
55
Addition polymerisation.
Addition polymers are made from alkenes
The names of polymers come from the name of the monomer

Propene makes poly(propene)

Chloroethene makes poly(chloroethene)

Ethene makes poly(ethene)
The monomer for poly(ethene) is ethene.
This is an unsaturated molecule containing a carbon to carbon double bond, which makes the
ethene molecule reactive.
The double bond splits open leaving free bonds to join onto other molecules forming a long
chain.
H H
C C
H H
Ethene
H H H H H H
H H
C C C C C C
C C
H H H H H H
H H
Poly(ethene)
Repeating unit
H
Butene can also be used to make a polymer,
called poly(butene.)
H H
H C C C C H
H H
To make things easier the molecule is drawn
in the shape of a letter H
56
H
CH3 H
CH3 H
CH3 H
CH3 H
CH3 H
C
C
C
C
C
C
C
C
H
CH3
H
CH3
H
CH3 H
Butene
Repeating unit
C
CH3 H
C
CH3
Poly(butene)
Since addition polymers are formed by opening a C=C and joining to give a chain they have
nothing but carbon atoms in the main chain.
57
Condensation polymerisation.
In condensation polymerisation many small molecules combine to form a polymer.
In condensation polymerisation molecules are joined together by removing
water.
 Polyesters are formed by condensation polymerisation from:
alcohols with two hydroxyl groups (–OH)
and carboxylic acids with two carboxyl groups (–COOH)
 This means that the polyester molecules can continue to grow in both directions
with many ester linkages.
In order to obtain the structures of the monomers used to make the polyester:
1. Locate the ester link:
58
O
C
O
2. Water adds on like this:
O
C
O
O
H
H
3. The carboxyl and hydroxyl groups in the monomers can now be formed:
O
C
O
H + H
O
…..
Natural Polymers
Many addition and condensation polymers are made by the chemical industry; however, polymers are found
in nature in both plants and animals.

Starch is a natural condensation polymer made from glucose in plants.

Proteins are also natural condensation polymers.

They are the major structural material of plant and animal tissue e.g. muscles, skin, hair and are
involved in the maintenance and regulation of life processes.
59
Unit 2
Metals
Metals conduct electricity because they have a sea of delocalised electrons.
LEO: Loss Electrons Oxidation
GER: Gain Electrons Reduction
Reactions of metals (OXIDATION)
1. Metals and water.
metal
e.g.
+ water
metal hydroxide
potassium + water
+ hydrogen
potassium hydroxide + hydrogen
(formula) 2K
+ 2H2O
2KOH +
H2
(ionic)
+ 2H2O
2K+
2OH- + H2
2K
+
2. Metals and acid.
Metals above hydrogen in the electrochemical series react with acids.
metal
e.g.
+ acid
magnesium + hydrochloric
acid
salt
+ hydrogen
magnesium +
chloride
hydrogen
(formula)
Mg + 2HCl
MgCl2
+ H2
(ionic)
Mg + 2H+ + 2Cl-
Mg2+ + 2Cl- + H2
3. Metals and oxygen
e.g.
metal
+
oxygen
metal oxide
magnesium
+
oxygen
magnesium oxide
(formula)
2Mg(s) +
O2(g)
2MgO(s)
(ionic)
2Mg(s)) +
O2(g)
2Mg2+O2- (s)
60
4. The Reactivity Series
These reactions give an indication of the reactivity of the metal and are summarised below. This is called
the reactivity series.
Metal
Reaction with
Oxygen
Potassium
Sodium
React
React
Lithium
Calcium
Water
react
Aluminium
with
form
forming
and hydrogen
Magnesium
Zinc
React
forming
metal hydroxide
to
Acid
steam
salt
and
Iron
Tin
metal oxide
hydrogen
Lead
Copper
No
Mercury
No
Silver
No
Gold
Reaction
Reaction
61
Reaction
Extracting metals from compounds (REDUCTION)
Metals are found as compounds called ores.
The more reactive metals form the most stable ores and so are hardest to obtain.
5. Methods of extraction.
Heating metal oxides
The least reactive metals can be obtained from their ores simply by heating
e.g.
Metal oxide
metal
+
oxygen
mercury oxide
mercury
+
oxygen
2HgO
2Hg
+
O2
Heating metal oxides with carbon
More reactive metals are extracted using carbon to remove the oxygen.
e.g.
Metal oxide
+
carbon
metal +
carbon dioxide
lead oxide
+
carbon
lead
+
carbon dioxide
PbO2
+
C
Pb
+
CO2
Electrolysis
Electricity is needed to obtain the most active metals from their compounds.
e.g.
aluminium oxide
aluminium
+
oxygen
4Al
+
3O2
(Al3+)2(O2-)3
4Al
+
3O2
Al3+ + 3e-
Al
2Al2O3
62
Reduction
More about Electrolysis
For electrolysis to happen the ionic lattice must be broken down by dissolving or melting the
compound.
This makes the ions free to move and so carry the current.
e.g. Electrolysis of molten lead (II) bromide- Pb2+(Br-)2 (l)
molten lead (II)
bromide
Product at the
negative electrode
lead
At the negative electrode : Pb
(metal ions gain electrons)
–
At the positive electrode : 2Br
(non-metal ions lose electrons)
2+
(aq) + 2e →
(aq) →
6. Percentage of Metal in an Ore.
Step 1 Calculate the gram formula mass (gfm)
K2CO3 - potassium carbonate
3 x 16 = 48
1 x 12 = 12
2 x 39 = 78
= 138g
Step 2 Calculate the total mass of the metal in the formula
K
3 x 39 = 78g
Step 3 Divide the mass of the metal by the gfm and multiply by 100%
% mass of the metal = 78 x 100%
138
= 56.5%
63
Product at the
positive electrode
bromine
Br2 (g) + 2e
Pb (s)
Making electricity
Electricity is a flow of charged particles:
flow of electrons through metals,
flow of ions through solutions or melts.
Electrons always flow from metals high in the electrochemical series through the wires to metals
lower in the electrochemical series.
The further apart two metals are in the electrochemical series the larger the voltage obtained.
.
2.7V
Mg
1.1V
Cu
Zn
Cu
electrolyte
64
Redox & Displacement
When electrons go into an electrode we get REDUCTION
When electrons leave an electrode we get OXIDATION
Electron flow
V
Zn
Ion Bridge
Zn2+SO42-
Cu
Cu2+SO42-
The copper ions in the solution accept these electrons and turn into copper metal.
Zn
Cu2+ + 2e-
Zn2+
Cu
+
2e-
oxidation
reduction
The oxidation and reduction equations can be combined to form what is known as a redox equation
e.g.
Zn + Cu2+
Zn2+
65
+ Cu
redox
Cells involving non-metals.
The electrochemical series in the data book has some reactions that involve non-metals
V
Zn
Ion Bridge
Zn2+SO42-
C
I2 / K+I-
If we are told that the following reaction occurs:
I2 + 2e-
2I-
This is Reduction which is where electrons go into an electrode.
So electrons must be going in at the C electrode.
This means that electrons must be leaving at the Zn electrode (Oxidation)
Zn2+
Zn
+
2e-
reduction
So electrons must be flowing from the Zinc to the carbon.
Electron flow
V
Zn
Ion Bridge
Zn2+SO42-
C
I2 / K+I-
The diagram shows that electrons flow from zinc metal to the carbon rod and the reactions are shown
below:
Zn
Zn2+ +
2ereduction
I2
+
2e
2I
oxidation
The redox equation will be:
I2
+
Zn
2I-
+
66
Zn2+
redox
Displacement
Metals can displace ions of a less reactive metal.
For example:
Zn(s) + Cu2+SO42-
Cu(s) + Zn2+SO42-
Omitting the sulphate ions (SO42-) and the state symbols makes the process clearer.
Zn
+ Cu2+
Cu
+ Zn2+
Being the more reactive metal the zinc metal is losing its electrons which it gives to the copper ions..
Zn
Zn2+ +
Cu2+ + 2e-
Cu
2 e-
oxidation
reduction
Adding the oxidation and reduction equations together gives us the redox equation
Zn
+ Cu2+
Cu
+ Zn2+
redox
Metals reacting with acids is an example of a displacement reaction.
Zn(s) + (H+)2SO42-
Zn2+SO42- + H2
The zinc metal is displacing the hydrogen ions.
The metals which do not react with acids, can not displace the hydrogen ions and so these metals
must be less reactive than hydrogen.
Therefore hydrogen is put into the reactivity series between lead and copper.
The process of losing electrons is called oxidation and that of gaining electrons is called reduction.
Fuel Cells & Rechargeable Batteries.
A fuel cell is a device that converts chemical energy from a fuel such as hydrogen into electrical
energy through a chemical reaction with oxygen or some other oxidising agent (a substance that
causes oxidation)
67
Fertilisers
Fertilisers contain the essential elements (NPK) that plants need.
Examples are
ammonium salts
(such as NH4Cl),
nitrates
(such as KNO3)
potassium salts
(such as K2SO4)
phosphates
(such as K3PO4)
Percentage composition
Fertilisers need to contain different proportions of N : P : K for different plants.
When comparing fertilisers it is useful to know the percentage of an element in that fertiliser.
calculate the percentage of nitrogen in ammonium nitrate we first calculate the GFM
NH4NO3
3x0
1xN
4xH
1XN
= 3 x 16 = 48
= 1 x 14 = 14
=4x1 = 4
= 1 X 14 = 14
GFM
= 80 g
From the total number of nitrogen atoms in ammonium nitrate calculate the total mass of nitrogen.
There are 2 nitrogen atoms in the formula therefore total mass nitrogen = 14 + 14 = 28
Percentage nitrogen = mass of nitrogen x 100
GFM
= 28 x 100
80
68
=35%
To
Making fertilisers
Making Fertilisers by Neutralisation
Nitrate fertilisers are formed using nitric acid:
HNO3
+
NaOH
NaNO3
+
H2O
Phosphate fertilisers are formed using phosphoric acid:
H3PO4
+
3NH3
(NH4)3PO4
For fertilisers containing potassium, an acid is reacted with a potassium compound:
K2CO3
+
HNO3
KNO3
+
H2O
+
CO2
Ammonium fertilisers are formed using ammonia and nitric acid:
NH3
+
HNO3
NH4NO3
The Haber Process
To make ammonium nitrate we need ammonia and nitric acid.
To make nitric acid we need ammonia so the production of ammonia is essential to making ammonium nitrate.
Ammonia is formed by reacting nitrogen (from air) and Hydrogen (from natural gas).
Iron is used as a catalyst and he reaction is carried out at 400oC.
N2(g) + H2(g)
NH3(g)
The reaction is reversible i.e. it can go in both directions.
To help stop this, the ammonia is cooled down to turn it into a liquid.
Any unreacted nitrogen and hydrogen is recycled.
Unreacted N2 and H2 recycled
Nitrogen
Hydrogen
Reaction
Chamber
Fe catalyst
400 0C
200 atm
N2/H2/NH3
Cooler
liquid NH3
The ammonia is then used to make nitric acid.
69
NUCLEAR CHEMISTRY
Types of radiation
There are 3 types of radiation, alpha ( ), beta (
an electrical field.
) and gamma ( ). Their properties can be studied using
α
γ
β
+
α
β
γ
paper
2cm Al
5cm concrete
Alpha particles
Alpha radiation consists of helium nuclei,
Alpha can only travel a few centimetres.
Beta particles
A beta particle is an electron,
.
Beta can travel metres.
Gamma waves
Gamma waves are electronegative waves.
They can travel 100’s of metres.
70
Nuclear equations
In nuclear equations, the sum of the atomic number and mass number on each side of the equation should
balance.
Alpha emissions
Beta emissions
Half-life
The half-life of a radioisotope never changes. It is the time taken for the sample's activity to fall by half.
Example 1
The mass of a radioisotope falls from 1.6g to 0.1g in 2 hours. What is the half-life of this
radioisotope?
Answer
1.6 g
0.8g
0.4g
0.2g
0.1g
The activity halves 4 times in 2 hours
Half life = 2/4 = 0.5 hour
Example 2
If a 1g sample of a radioisotope with a half-life of 3 days has an activity of 32 c.p.m., how long
would it take for the activity to fall to 8 c.p.m.?
Answer If 1g sample is 32 cpm.
32 cpm
16 cpm
Activity halves 2 times.
Half life = 3 days
Time taken = 2 x 3 = 6 days
8 cpm
71
Carbon dating
Carbon -14 is present in the atmosphere.
Carbon dioxide is responsible for carbon-14 entering the food chain.
The levels of carbon dioxide in living things stays constant as it is absorbed through photosynthesis or in
food, and it decays by beta emission.
When living things die no more carbon-14 is absorbed and as the C-14 continue to decay by beta emission
its levels fall.
We can use the half-life of C-14 (5,700 years) to calculate the age of the object.
Example
If a piece of wood has 25 % the radiocativity of a living tree then it has undergone 2 half-lives.
100%
50%
25%
Half-life = 5700 years
Tree died 2 x 5,700 = 11,400 years ago.
Uses of radioisotopes
Radioisotopes of elements have a wide variety of uses.
Cobalt-60
Used in medicine to treat cancer..
Iodine-133
Used to treat thyroid gland.
72