Stability of Noble Gases
[edit] What is Chemical Compound
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A chemical compound is a substance that is formed by more than one elements
that bond together chemically in a fixed proportions.
In periodic table, there are only 118 elements, and about 1/3 of them are synthetic
elements.
Only a few substances exist as element (Not Compound) in nature.
The table below shows some examples of substance exist as element in nature.
Element exist as
monoatomic gas
Element exist as
diatomic Molecule(gas)
Element exist as
solid
Helium (He)
Oxygen (O2)
Carbon (Graphite
& Diamond)
Neon (Ne)
Argon (Ar)
Krypton (Kr)
Xenon (Xe)
Radon (Rn)
Nitrogen (N2)
Gold
Silver
Platinum
In nature, we can find millions of substances, which means most of the chemical
substances exist as compound in nature.
In short, elements tend to form compound in nature.
Why Elements Tend to Form Compound?
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A compound is formed by 2 or more elements hold together by a force called
chemical bond.
Before studying why elements like to bond together, we need to know why certain
elements such as Helium and Neon do not form any bonds with other elements.
Why Noble Gases Don't Form Compound
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In previous chapter, we have discussed that Group 18 elements (Noble Gases)
exist as monoatom in nature.
They are inert in nature and do not react with any other elements (or themselves)
to form any chemical compounds.
In other words, they are chemically very very stable (or chemically very very nonreactive).
[edit] Duplet and Octet Electron Arrangement
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The charge on the nucleus and the number of electrons in the valence shell
determine the chemical properties of an atom.
The stability of noble gas is due to their electrons arrangement.
The diagram above shows the first four elements of Noble Gas.
We can see that the outer most shell (valence shell) of Helium has 2 electrons. We
call this duplet electron arrangement. We should take notes that the maximum
number of electrons can be filled in the first shell is 2 electrons, which means 2
electrons in the first shell is considered FULL.
The valence shell all other Group 18 elements (including Xenon and Radon which
is not shown in the diagram) has 8 electrons, and we call this octet electron
arrangement.
When the electron arrangement of an atom is duplet or octet, the energy of the
electrons is very low, and it is very difficult (even though it is not impossible) to
add or remove electrons from the atom.
This explain why noble gases are reluctant to react with all other elements.
[edit] The Octet Rule
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So far we have learnt that the electon arrangement of noble gases are octet duplet,
and this is the most stable electron arrangement of an atom.
Atoms of other main group elements which is not octet tend to react with other
atoms in various ways to achieve the octet.
The tendency of an atom to achieve an octet arrangement of electrons in the
outermost shell is called the octet rule.
If the outermost shell is the first shell, then the maximum number of electrons is
two, and the most stable electron arrangement will be duplet.
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A configuration of two electrons in the first shell, with no other shells occupied
by electrons, is as stable as the octet electron arrangement and therefore is also
said to obey the octet rule.
Important Notes
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Most of the elements (except
noble gases) are chemically not
stable.
It is the aim of every atom to
achieve the duplet or octet
electron arrangement. This makes
them very stable.
It is only the valence electrons in
the outermost shell involved in
bonding. The electrons in the
inner shells are not involved.
The maximum number of
electrons in the first shell is two.
This is called a duplet.
The maximum number in the
second shell is eight. This is
called an octet.
[edit] How Atoms Achieve Duplet or Octet Electron
Arrangement?
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Atoms can achieve duplet or octet electron arrangement in 3 ways:
1. throw away the excess electron(s)
2. receiving electron(s) form other atom if they are lack of electron(s)
3. sharing electron
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2 types of chemical bonds are commonly formed between atoms, namely
1. Ionic Bond
2. Covalent Bond
[edit] The Ionic Bond
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By releasing or receiving electron(s), the atoms will become ions and
consequently form ionic bond between the ions.
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Ionic bonds are always form between metal and non-metal. For example, sodium
(metal) react with chlorine (non-metal) will form an ionic bond between sodium
ion and chloride ion.
The compounds formed is called the ionic compound.
Some time, an ionic bond is also called electrovalent bond.
[edit] The Covalent Bond
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By sharing electron(s), the atoms will form covalent bond between the atom and
the molecule formed is call the covalent molecule.
Covalent bond is always formed between non-metal with another non-metal.
5.2 Formation of Ion
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An ion is an atom or group of atoms carrying positive or negative charge.
Example Ca2+, O2-, SO42- etc.
If a particle has equal numbers of protons (+) and electrons (-), the particle charge
is zero, and the particle is said to be neutral.
In a chemical reaction, electron(s) can be transfered from atom to another atom.
If electron(s) is removed from an atom, the number of protons will be more than
number of electrons. In this case, the atom will has excess positive charge and
hence form a positive ion (cation).
If aton gains negative electrons, there is an excess negative charge in the atom, so
a negative ion is formed.
In other words,
1. The atom losing electrons forms a positive ion (cation) and is usually a metal.
2. The atom gaining electrons forms a negative ion (anion) and is usually a nonmetallic element.
Formation of Negative Ion
Example: Formation of Fluoride Ion
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A fluorine atom has 9 protons and 9 electrons.
Since the number of protons is equals to the number of electrons, the fluorine
atom is neutral.
The electron arrangement of fluorine atom is 2.7. This is not a stable arrangement
of electrons.
To achieve the stable electron arrangement of noble gases (octet electron
arrangement), fluorine need to receive 1 electron from the other atom.
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As shown in the diagram above, after receiving 1 electron, the electron
arrangement of fluorine become 2.8, which is an octet arrangement of electrons.
At the same time, the number of electrons has increased by 1 and become 10
electrons while the number of proton remain unchanged.
Hence the charge of the fluoride ion is -1.
The table below shows the difference between a fluorine atom and a fluoride ion
in term of its number of proton and electron and its electron arrangement.
Fluorine Atom
Number of proton = 9
Number of electron = 9
Eelctron Arrangement = 2.7 (Not Octet)
Charge = 0
Fluoride Ion
Number of proton = 9
Number of electron = 10
Eelctron Arrangement = 2.8 (Not Octet)
Charge = -1
[edit] Difference Between Fluoride Ion and Neon Atom
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After forming an ion, the electron arrangement of the ion is similar to the noble
gases.
For example, the electron arrangement of fluoride ion is 2.8 which is similar to a
Neon atom, 2.8.
By referring to the electron arrangement, sometime, students may mistaken an ion
as noble gas.
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We should take notes that even though the fluoride ion and the neon atom have
similar electron arrangement, fluoride ion carry charge whereas neon atom is
neutral.
[edit] Formation of Positive Ion
Example: Formation of Magnesium Ion
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A magnesium atom has 12 protons and 12 electrons and it is neutral.
The electron arrangement of magnesium atom is 2.8.2 .
To achieve the stable electron arrangement of noble gases (octet electron
arrangement), magnesium atom need to loss2 electrons.
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As shown in the diagram above, after losing 2 electrons, the electron arrangement
of magnesium become 2.8, which is an octet arrangement of electrons.
At the same time, the number of electrons has reduced by 2 and become 10
electrons while the number of proton remain unchanged.
Hence the charge of the magnesium ion is +2.
The table below shows the difference between a magnesium atom and a
magnesium ion in term of its number of proton and electron and its electron
arrangement.
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Fluorine Atom
Fluoride Ion
Number of proton = 9
Number of electron = 9
Eelctron Arrangement = 2.7 (Not Octet)
Charge = 0
Number of proton = 9
Number of electron = 10
Eelctron Arrangement = 2.8 (Not Octet)
Charge = -1
Ionic Bonding 1. Ionic bonds are formed by one atom transferring electrons to another
atom to form ions. Ions are atoms, or groups of atoms, which have lost or gained
electrons. 2. The atom losing electrons forms a positive ion (a cation) and is usually a
metal. The overall charge on the ion is positive due to excess positive nuclear charge
(protons do NOT change in chemical reactions). 3. The atom gaining electrons forms a
negative ion (an anion) and is usually a non-metallic element. The overall charge on the
ion is negative because of the gain, and therefore excess, of negative electrons. 4. Ions of
opposite charge will attract one another, thus creating an ionic bond. 5. The examples
below combining a metal from Groups 1 (Alkali Metals), or 2, with a non-metal from
Group 6 or Group 7 (The Halogens)
[edit] Example
[edit] A Group 1 metal + a Group 7 non-metal
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In terms of
electron
arrangement, the
sodium donates its
outer electron to a
chlorine atom
forming a single
positive sodium
ion and a single
negative chloride
ion.
The valencies of
Na and Cl are both
1, that is, the
numerical charge
on the ions. NaF,
KBr, LiI etc. will
all be
electronically
similar.
The atoms have
become stable
ions, because
electronically,
sodium becomes
like neon and
chlorine like
argon.
[edit] A Group 2 metal + a Group 7 non-metal
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In terms of
electron
arrangement, the
magnesium
donates its two
outer electrons to
two chlorine atoms
forming a double
positive
magnesium ion
and two single
negative chloride
ions.
The atoms have
become stable
ions, because
electronically,
magnesium
becomes like neon
and chlorine like
argon.
NOTE you can
draw two separate
chloride ions, but
in these examples
a number subscript
has been used, as
in ordinary
chemical formula.
The valency of Mg
is 2 and chlorine 1,
ie the numerical
charges of the
ions. BeF2,
MgBr2, CaCl2 or
CaI2 etc. will all
be electronically
similar.
[edit] A Group 2 metal + a Group 6 non-metal
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In terms of
electron
arrangement, one
Magnesium atom
donates its two
outer electrons to
one oxygen atom.
This results in a
double positive
magnesium ion to
one double
negative oxide ion.
All the ions have
the stable
electronic
structures 2.8.8
(argon like) or 2.8
(neon like). the
valency of both
calcium and
oxygen is 2.
MgO, MgS, or CaS will
be similar electronically
(S and O both in Group 6)
[edit] Predicting The Formula for Ionic Compound
Formula of the ionic
compound
Element that combine
Element X
from
Charge of the
ion
Element Y
from
Charge of the
ion
Group I
+1
Group V
-3
Group I
+1
Group VI
-2
Group I
+1
Group VII
-1
Group II
+2
Group V
-3
Group II
+2
Group VI
-2
Group II
+2
Group VII
-1
Group III
+3
Group V
-3
Group III
+3
Group VI
-2
Group III
+3
Group VII
-1
5.3 Covalent Bonding
1. Covalent bonds are formed by atoms sharing electrons to form molecules. This type of
bond usually formed between two non-metallic elements. 2. The molecules might be that
of an element ie one type of atom only OR from different elements chemically combined
to form a compound. 3. The covalent bonding is caused by the mutual electrical attraction
between the two positive nuclei of the two atoms of the bond, and the electrons between
them. 4. One single covalent bond is a sharing of 1 pair of electrons, two pairs of shared
electrons between the same two atoms gives a double bond and it is possible for two
atoms to share 3 pairs of electrons and give a triple bond.
Examples of Covalent Compound
Chlorine phosphorus oxygen carbon dioxide nitrogen Tetrachloro-methane Hydrogen
sulfur dioxide
[edit] Number of Bond
[edit] Example of Covalent Bonding
[edit] Single bond
[edit] Hydrogen
[edit] Fluorine
[edit] Water
[edit] Ammonia
[edit] Tetrachloromethane
[edit] Double Bond
[edit] Oxygen
[edit] Carbon Dioxide
Image:FormationCO2.png
[edit] Triple Bond
[edit] Nitrogen
Number of Covalent bonds
depends on how many pairs of electron been shared.
[edit] Predicting The Molecular Formula
Element that combine
Formula of the ionic compound
Element X from Valency Element Y from Valency
Group V
3
Group V
3
Group V
3
Group VI
2
Group V
3
Group VII
1
Group VI
2
Group VI
2
Group VI
2
Group VII
1
Group VII
1
Group VII
1
5.4 Properties of Ionic Compounds
[edit] The Structure of Ionic Compounds--Crystal Lattices
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The alternate positive and negative ions in an ionic solid are arranged in an
orderly way in a giant ionic lattice structure shown on the left.
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The ionic bond is the strong electrical attraction between the positive and negative
ions next to each other in the lattice.
The bonding extends throughout the crystal in all directions.
Salts and metal oxides are typical ionic compounds.
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Some of these compounds, like magnesia (MgO) and alumina (Al2O3), are so
stable that they are used as refractory material, to line the inside of furnaces. Such
substances must be stable up to at least 1500 °C.
Another property of crystal lattices is that they are non-conductors of electricity.
This is because the ions are in fixed positions and are unable to move.
[edit] Properties of Ionic Compounds
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This strong bonding force makes the structure hard (if brittle) and has high
melting and boiling points, so they are not very volatile!
The bigger the charges on the ions the stronger the bonding attraction eg
magnesium oxide Mg2+O2- has a higher melting point than sodium chloride
Na+Cl-.
Unlike covalent molecules, ALL ionic compounds are crystalline solids at room
temperature.
They are hard but brittle, when stressed the bonds are broken along planes of ions
which shear away. They are NOT malleable like metal.
Many ionic compounds are soluble in water but not all, so don't make
assumptions.
The solid crystals DO NOT conduct electricity because the ions are not free to
move to carry an electric current.
However, if the ionic compound is melted or dissolved in water, the liquid will
now conduct electricity, as the ion particles are now free.
[edit] Properties of Covalent Compounds
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Covalent compounds can be divided into those which form small (simple)
independent molecules and those which form giant molecular lattices.
[edit] Simple Molecule
[edit] Structure
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These are made up of independent molecular units, as shown in Figure 6.7.
As there are no ions formed, the attractive forces between molecules in solid,
covalent compounds like iodine and sulphur are much weaker.
They are called van der Waals' forces and produce a weak, molecular lattice with
low melting points.
n covalent liquids like water, the molecules are even further apart, so the van der
Waals' forces are weaker still, and in covalent gases like ammonia and methane,
these forces are almost non-existent.
However, in water, there are other attractive forces between molecules. These
forces are called hydrogen bonds and they give water much higher melting and
boiling points than expected with such weak van der Waals' forces.
[edit] Properties Of Simple Covalent Molecular Substances - Small
Molecules!
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The electrical forces of attraction, that is the chemical bond*, between atoms in
any molecule are strong and most molecules do not change chemically on
moderate heating.(* sometimes referred to as the intramolecular bond)
However, the electrical forces** between molecules are weak and easily
weakened further on heating.
These weak attractions are known as **intermolecular forces and consequently
the bulk material is not usually very strong.
Consequently small covalent molecules tend to be volatile liquids, easily
vapourised, or low melting point solids.
On heating the inter-molecular forces are easily overcome with the increased
kinetic energy gain of the particles and so have low melting and boiling points.
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They are also poor conductors of electricity because there are no free electrons or
ions in any state to carry electric charge.
Most small molecules will dissolve in a solvent to form a solution.
[edit] Macromolecular compounds
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These have giant, covalent molecules with extremely large molecular lattices.
They are very stable, as all the atoms are joined together by strong covalent bonds
to give a giant three-dimensional lattice.
Often the lattice is tetrahedral in shape, as every atom is covalently linked to four
others.
Examples of such macromolecules are diamond and sand (see Figure 6.8).
[edit] Diamond and Silica(Sand)
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A diamond crystal or a grain of sand is just one giant molecule. Such molecules,
because they are so rigid and strong, have very high melting points.
Large Covalent Molecules And Their Properties
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This type of structure is thermally very stable and they have high melting and
boiling points.
They are usually poor conductors of electricity because the electrons are not
usually free to move as they can in metallic structures.
Also because of the strength of the bonding in all directions in the structure, they
are often very hard, strong and will not dissolve in solvents like water.
Silicon dioxide (silica, SiO2) has a similar 3D structure and properties, shown
below diamond.
The hardness of diamond enables it to be used as the 'leading edge' on cutting
tools.
[edit] Graphite
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Diamond is an allotrope of carbon. Allotropes are different forms of the same
element in the same physical state
Oxygen O2 (dioxygen) and ozone O3 (trioxygen) are two gaseous allotropes of
the element oxygen.
Carbon also occurs in the form of graphite. The carbon atoms form joined
hexagonal rings forming layers 1 atom thick.
There are three strong covalent bonds per carbon (3 C-C bonds in a planar
arrangement from 3 of its 4 outer electrons), BUT, the fourth outer electron is
'delocalised' or shared between the carbon atoms to form the equivalent of a 4th
bond per carbon atom.
The layers are only held together by weak intermolecular forces shown by the
dotted lines NOT by strong covalent bonds.
Like diamond and silica (above) the large molecules of the layer ensure graphite
has typically very high melting point because of the strong 2D bonding network
(note: NOT 3D network)..
Graphite will not dissolve in solvents because of the strong bonding but there are
two crucial differences compared to diamond ...
Electrons, from the 'shared bond', can move freely through each layer, so graphite
is a conductor like a metal (diamond is an electrical insulator and a poor heat
conductor). Graphite is used in electrical contacts eg electrodes in electrolysis.
The weak forces enable the layers to slip over each other so where as diamond is
hard material graphite is a 'soft' crystal, it feels slippery. Graphite is used as a
lubricant.
These two different characteristics described above are put to a common use with
the electrical contacts in electric motors and dynamos. These contacts (called
brushes) are made of graphite sprung onto the spinning brass contacts of the
armature. The graphite brushes provide good electrical contact and are selflubricating as the carbon layers slide over each other.